Download Chapter 20

Chapter 20
Oxidation-Reduction Reactions
Some Common Reactions
The combustion of gasoline in an automobile
engine requires oxygen
Burning of wood in a fireplace requires oxygen
The reactions that break down food in your
body and release energy use oxygen
The oxide of hydrogen is water
Charcoal oxidizes when it burn forming CO2
Bleaching stains in fabric is still oxidation even
though it does not burn.
Oxidation
When methane burns in air, it oxidizes and
forms oxides of carbon and hydrogen.
CH4(g) + 2O2(g)
CO2(g) + 2H2O(g)
When elemental iron turns to rust, it slowly
oxidizes to compounds such as iron (III)
oxide.
4Fe(s) + 3O2(g)
2Fe2O3(s)
Reduction
Originally reduction meant a loss of oxygen
from a compound
2Fe2O3(s) + 3C(s)
iron oxide
carbon
4Fe(s) + 3CO2(g)
iron
carbon dioxide
Reduction of iron ore to metallic iron involves
the removal of oxygen from iron (III) oxide.
Involves heating the ore with carbon.
Question
What happens to magnesium and oxygen
when they react to form magnesium oxide?
+2
2Mg
magnesium
+
O2
oxygen
-2
2MgO
magnesium oxide
Magnesium loses electrons to form Mg2+
Oxygen gains electrons to form O2-
Electron Shift in Redox Reactions
The modern concept of oxidation and reduction
have been extended to include many
reactions that do not even involve oxygen.
Oxygen is the most electronegative element
(besides fluorine)
When oxygen bonds with an atom of a different
element (except fluorine), electrons from that
atom shift toward oxygen.
Redox Reactions
Redox reactions are currently understood to
involve any shift of electron between reactants.
Oxidation – a process that involves a complete or
partial loss of electrons or a gain of oxygen.
• Results in an increase in the oxidation number of an atom
Reduction – a process that involves a complete or
partial gain of electrons or the loss of oxygen.
• Results in a decrease in the oxidation number of an atom
Redox Reactions
Oxidation and reduction always occur
simultaneously.
The substance gaining oxygen or losing
electrons is oxidized
The substance losing oxygen or gaining
electrons is reduced.
Redox Reactions
During a reaction between a metal and a nonmetal,
electrons are transferred from atoms of the metal
to atoms of the nonmetal.
Mg +
magnesium
S
sulfur
Mg2+S2magnesium sulfide
2 electrons are transferred from a magnesium atom
to a sulfur atom.
Magnesium atoms become more stable by the loss
of electrons. Sulfur atoms become more stable
by the gain of electrons
Redox Reactions
Mg +
magnesium
Oxidation: Mg
S
sulfur
Mg2+S2magnesium sulfide
Mg2+ + 2e- (loss of electrons)
Reduction: S + 2e-
S2- (gain of electrons)
Magnesium atom is said to be oxidized to a
magnesium ion
Sulfur atom is said to be reduced to a sulfide ion.
Redox Reactions
When a metal combines with oxygen, it loses
electrons
When oxygen is removed from the oxide of a metal,
the metal gains electrons.
This knowledge is what led to the broader definition
of oxidation and reduction as an exchange of
electrons.
Redox Reactions
Reducing agent – the substance that loses the
electrons
Mg
+
S
MgS
reducing agent
oxidized
Oxidizing agent – the substance that accepts
electrons
Mg + S
MgS
oxidizing agent
reduced
Sample Problem
Silver nitrate reacts with copper to form copper nitrate
and silver. From the equation below, determine what
is oxidized and what is reduced. Identify the oxidizing
agent and the reducing agent.
2AgNO3 + Cu
Cu(NO3)2 + 2Ag
Rewrite the equation in ionic form
2Ag+ + 2NO3- + Cu
Cu2+ + 2NO3- + 2Ag
Sample Problem
2Ag+ + 2NO3- + Cu
Cu2+ + 2NO3- + 2Ag
Oxidation: 2 e- are lost from copper when it
becomes a Cu2+
Reduction: 2e- are gained by two silver ions which
become neutral silver atoms.
2Ag+ + 2NO3- + Cu
oxidizing
Agent
reduced
reducing
agent
oxidized
Cu2+ + 2NO3- + 2Ag
Sample Problem
2Na
+
oxidized
reducing agent
4Al
+
Oxidized
Reducing agent
2Na+
S
+
S2-
reduced
oxidizing agent
4Al3+
3O2
reduced
oxidizing agent
2I-
I2 + 2eoxidation
Zn2+
+
2e-
reduction
Zn
+
3O2-
Redox with Covalent Compounds
Some reactions involve covalent compounds. In
these compounds complete electron transfer
does not occur.
2H2 (g) + O2 (g)
2H2O (l)
In each reactant hydrogen molecule, the bonding
electrons are shared equally between the
hydrogen atoms.
In water, however, the bonding electrons are
pulled toward oxygen because it is much more
electronegative than hydrogen.
Redox with Covalent Compounds
2H2 (g) + O2 (g)
2H2O (l)
There is a shift of bonding electrons away from
hydrogen, even though there is not a complete
transfer.
Hydrogen is oxidized because it undergoes a
partial loss of electrons.
Redox with Covalent Compounds
2H2 (g) + O2 (g)
2H2O (l)
In oxygen, the bonding electrons are share
equally between oxygen atoms in the reactant
oxygen molecule.
When oxygen bonds to hydrogen in the water
molecule, there is a shift of electrons toward
oxygen.
Oxygen is thus reduced because it undergoes a
partial gain of electrons.
Redox with Covalent Compounds
H
H
O
O
e- shared
e- shared
equally
equally
H is reducing
agent
O is oxidizing
agent
H
O
shift of bonding
e- away from H
H
and toward O
Redox with Covalent Compounds
Some reactions involving covalent reactants or
products, the partial electron shifts are less
obvious.
General guideline for covalent reactants or
products:
• for carbon compounds, the addition of oxygen
or the removal of hydrogen is always oxidation
Processes Leading to
Oxidation & Reduction
Processes Leading to Oxidation & Reduction
Oxidation
Reduction
Complete loss of electrons
(ionic reactions)
Complete gain of electrons
(ionic reactions)
Shift of electrons away from
an atom in a covalent bond
Shift of electrons toward an
atom in a covalent bond
Gain of oxygen
Loss of oxygen
Loss of hydrogen by a
covalent compound
Gain of hydrogen by a
covalent compound
Increase in oxidation
number
Decrease in oxidation
number
Corrosion
Iron, often used in the
form of the alloy steel,
corrodes by being
oxidized to ions of iron
by oxygen.
2Fe(s) + O2 (g) + 2H2O (l)
4Fe(OH)2(s) + O2 (g) + 2H2O (l)
2Fe(OH)2 (s)
4Fe(OH)3 (s)
• Equations describe the corrosion of iron to iron
hydroxides in moist conditions.
Corrosion
Water in the environment accelerates the rate
of corrosion.
Corrosion occurs more rapidly in the presence of
salts and acids.
Salts and acids produce electrically
conducting solutions that make electron
transfer easier.
Resistance to Corrosion
Not all metals corrode easily.
Gold and platinum are called noble metals
because they are very resistant to losing their eby corrosion.
Other metals lose electrons easily but are
protected from extensive corrosion by the
oxide coating formed on their surface.
• Aluminum oxidizes quickly in air to form a
coating of very tightly packed aluminum oxide
particles.
Resistance to Corrosion
Iron also forms a coating when it corrodes
But the coating of iron oxide that forms is not
tightly packed.
Water and air can penetrate the coating and
attack the iron metal below it.
Corrosion continues until the iron object becomes
only a pile of rust.
Controlling Corrosion
To prevent corrosion, the metal surface can be
coated with oil, paint, plastic or another
metal.
Coatings exclude air and water from the
surface, preventing corrosion.
If coating is scratched or worn away, however, the
exposed metal will begin to corrode.
Controlling Corrosion
Another method of corrosion control
One metal is “sacrificed” or allowed to corrode, in
order to save a second metal.
To protect an iron object, a piece of
magnesium (or other active metal) may be
placed in electrical contact with the iron.
When oxygen and water attack the iron object, the
iron atoms lose electrons as the iron being to be
oxidized.
Controlling Corrosion
Because magnesium is a better reducing agent
than iron and is more easily oxidized
the magnesium immediately transfers electrons
to the iron, preventing their oxidation to iron
ions.
Magnesium is “sacrificed” by oxidation and
protects the iron in the process.
Controlling Corrosion
Sacrificial zinc and magnesium blocks are sometimes
attached to piers and ship hulls to prevent corrosion
damage in areas submerged in water.
Underground pipelines and storage tanks may be
connected to magnesium block for protection
It is easier and cheaper to replace a block of
magnesium or zinc than to replace a bride or a
pipeline.
Question
Can you identify the common chemical
characteristic of all metal corrosion?
The transfer of electrons from metals to oxidizing
agents.
Questions
Define oxidation and reduction in terms of the
gain or loss of oxygen.
Oxidation is the gain of oxygen
Reduction is the loss of oxygen
Define oxidation and reduction in terms of the
gain or loss of electrons.
LEO the lion goes GER
Loss of Electrons is Oxidation
Gain of Electrons is Reduction
Questions
What happens to the atoms in an iron nail that
corrodes?
Iron atoms are oxidized when iron corrodes
How do you identify the oxidizing agent and
the reducing agent in a redox reaction?
The species reduced is the oxidizing agent.
The species oxidized is the reducing agent.
Questions
Use electron transfer or electron shift to identify what
is oxidized and what is reduced in each reaction.
(use electronegativity values for molecular
compounds)
2Na(s) + Br2(l)
2NaBr(s)
Na oxidized, Br2 reduced
H2(g) + Cl2(g)
2HCl(g)
H2 oxidized, Cl2 reduced
Questions
Use electron transfer or electron shift to identify what
is oxidized and what is reduced in each reaction.
(use electronegativity values for molecular
compounds)
2Li(s) + F2(g)
2LiF(s)
Li oxidized, F2 reduced
S(s) + Cl2(g)
SCl2(g)
S oxidized, Cl2 reduced
Questions
Use electron transfer or electron shift to identify what
is oxidized and what is reduced in each reaction.
(use electronegativity values for molecular
compounds)
N2(g) + 2O2(g)
2NO2(s)
N2 oxidized, O2 reduced
Mg(s) + Cu(NO3)2(aq)
Mg(NO3)2(aq) + Cu(s)
Mg oxidized, Cu reduced
Questions
Identify the reducing agent and the oxidizing agent
for each reaction.
2Na(s) + Br2(l)
2NaBr(s)
Na reducing agent, Br2 oxidizing agent
H2(g) + Cl2(g)
2HCl(g)
H2 reducing agent, Cl2 oxidizing agent
Questions
Identify the reducing agent and the oxidizing agent
for each reaction.
2Li(s) + F2(g)
2LiF(s)
Li reducing agent, F2 oxidizing agent
S(s) + Cl2(g)
SCl2(g)
S reducing agent, Cl2 oxidizing agent
Questions
Identify the reducing agent and the oxidizing agent
for each reaction.
N2(g) + 2O2(g)
2NO2(s)
N2 reducing agent, O2 oxidizing agent
Mg(s) + Cu(NO3)2(aq)
Mg(NO3)2(aq) + Cu(s)
Mg reducing agent, Cu oxidizing agent
End of Section 20.1
Oxidation Numbers
Oxidation number is a + or – number assigned to
an atom to indicate its degree of oxidation or
reduction.
General Rule
A bonded atom’s oxidation # is the charge that it
would have if the e- in the bond were assigned to
the atom of the more electronegative element.
Rules for Assigning Oxidation Numbers
1. The oxidation number of a monatomic ion is = in
magnitude and sign to its ionic charge.
Bromide (Br1-) is -1
Iron III (Fe3+) is +3
2. The oxidation number of hydrogen in a
compound is +1, except in metal hydrides, such
as NaH, where it is -1
3. The oxidation number of oxygen in a compound
is -2 except in peroxides, such as H2O2, where it
is -1 and in compounds with the more
electronegative fluorine, where it is positive.
Rules for Assigning Oxidation Numbers
4. The oxidation number of an atom in uncombined
(elemental) form is 0.
Potassium metal (K) is 0
Nitrogen Gas (N2) is 0
5. For any neutral compound, the sum of the
oxidation numbers of the atoms in the compound
must equal 0
6. For polyatomic ion, the sum of the oxidation
numbers must equal the ionic charge of the ion.
Some Thought
Determining oxidation numbers of elements in
compounds is a way for chemists to keep track of
electron transfer during redox reactions
What are other examples where items are
numbered to keep track of movement?
The numbers on a sports player’s jersey
The area codes assigned to telephone numbers in
different regions
Binary Ionic Compounds
In binary ionic compounds, such as NaCl and
CaCl2, the oxidation numbers of the atoms equal
their ionic charges.
Na1+ + Cl-1
NaCl
oxidation # +1 -1
neutral
Ca2+ + Cl-1
CaCl2
oxidation # +2 -1
neutral
Note the sign I put before the oxidation number
Molecular Compounds
No ionic charges are associated with atoms of
molecular compounds.
However, oxygen is reduced in the formation of
water for example.
In water the two shared e- in the H – O bond are
shifted toward the O and away from the H.
Imagine the e- contributed by the two H atoms are
completely transferred to the O.
Molecular Compounds
The charge that would result from the transfer are
the oxidation numbers of the bonded elements.
The oxidation number of O is -2 and the oxidation
number of each hydrogen is +1
Oxidation numbers are often written above the
chemical symbols in a formula.
+1 -2
H2O
Multiple Oxidation Numbers
Many elements can have several different oxidation
numbers.
+1 +6 -2
K2CrO4 – Potassium Chromate
+1 +12 -2
K2CrO7 – Potassium Dichromate
Sample Problems
What is the oxidation number of each kind of atom
in the following ions and compounds?
+4 -2
SO2
+4 -2
CO32+1 +6 -2
Na2SO4
-3 +1
-2
(NH4)2S
Sample Problems
Determine the oxidation number of each element in
the following.
+3 -2
S2O3
+1
-1
Na2O2
+5
-2
P2O5
+5
-2
NO3-
Sample Problems
Determine the oxidation number of chlorine in each
of the following substances.
+1 +5 -2
KClO3
0
Cl2
+2
+7
-2
Ca(ClO4)2
+1
-2
Cl2O
Sample Problems
What are the oxidation numbers of iodine in the
following?
+1 +7 -2
HIO4
+1 +5 -2
HIO3
+1 +1 -2
HIO
0
I2
+1
-1
HI
Oxidation Number Changes
When copper wire is placed in a solution of silver
nitrate the following reaction occurs
+1 +5 -3
0
2AgNO3(aq) + Cu(s)
+2
+5 -2
0
Cu(NO3)2(aq) + 2Ag(s)
Ag is reduced from +1 to 0
Cu is oxidized from 0 to +2
Oxidation Number Changes
An increase in the oxidation number of an atom or
ion indicates oxidation.
A decrease in the oxidation number of an atom or
ion indicates reduction.
+2 +6 -2
0
2CuSO4(aq) + Fe2(s)
0
+2 +6 -2
2Cu(s) + 2FeSO4
Cu is reduced from +2 to 0
Fe is oxidized from 0 to +2
Sample Problem
Identify which atoms are oxidized and which are
reduced in the following reaction. Also identify
the oxidizing agent and the reducing agent.
+1 -1
0
2HBr(aq) + Cl2(g)
+1 -1
0
2HCl(aq) + Br2(l)
Cl is reduced from 0 to -1, so Cl2 is the oxidizing agent
Br is oxidized from -1 to 0, so Br1- is the reducing agent
Sample Problem
Identify which atoms are oxidized and which are
reduced in each reaction
0
0
O2(g) + 2H2(g)
+1 -2
2H2O(l)
O2 is reduced from 0 to -2
H2 is oxidized from 0 to +1
Sample Problem
Identify which atoms are oxidized and which are
reduced in each reaction
+1 +5 -2
+1 +3 -2
0
2KNO3(s)
2KNO2(s) + O2(g)
N is reduced from +5 to +3
O is oxidized from -2 to 0
Sample Problem
Identify the oxidizing agent and the reducing agent
in each equation.
0
0
O2(g) + 2H2(g)
+1 -2
2H2O(l)
O2 is reduced, thus O2 is the oxidizing agent
H2 is oxidized, thus H2 is the reducing agent
Sample Problem
Identify the oxidizing agent and the reducing agent
in each equation.
+1 +5 -2
+1 +3 -2
0
2KNO3(s)
2KNO2(s) + O2(g)
N is reduced, thus N is the oxidizing agent
O is oxidized, thus O is the reducing agent
Questions
What is the general rule for assigning oxidation
numbers?
The oxidation number is the charge a bonded atom
would have if the electrons in the bond were
assigned to the more electronegative element
How is a change I oxidation number related to the
process of oxidation and reduction?
An increase in oxidation number indicates oxidation; a
decrease in oxidation number indicates reduction.
Sample Problem
Identify which atoms are oxidized and which are
reduced in each reaction. Also identify the
oxidizing agent and the reducing agent.
0
0
2Na(s) + Cl2(g)
+1 -2
2NaCl(s)
Cl2 is reduced, thus Cl2 is the oxidizing agent
Na is oxidized, thus Na is the reducing agent
Sample Problem
Identify which atoms are oxidized and which are
reduced in each reaction. Also identify the
oxidizing agent and the reducing agent.
0
0
2HNO3(aq) + 6HI(aq)
+1 -2
2NO(g) + 3I2(s) + 4H2O(l)
O2 is reduced, thus O2 is the oxidizing agent
H2 is oxidized, thus H2 is the reducing agent
End of section 20.2
Identifying Redox Reactions
In general, all chemical reaction can be
assigned to one of two classes
1. Redox reactions in which electrons are
transferred from one reacting species to another.
a. Many single-replacement reactions,
combination reactions, decomposition
reactions and combustion reactions are redox
reactions.
2. All other reactions in which no electron transfer
occurs.
a. Double-replacement reactions and acid-base
reactions are not redox reactions
What Kind of Reaction Is It?
2Mg(s) + O2(g)
2MgO(s)
Combination reaction – two or more substances
react to form a single new substance
2HgO (s)
2Hg (l) + O2 (g)
Decomposition reaction – a single compound
breaks down into two or more simpler products.
Many are redox reactions
What Kind of Reaction Is It?
Zn(s) + Cu(NO3)2(aq)
Cu(s) + Zn(NO3)2 (aq)
Single Replacement Reaction – one element
replaces a second element in a compound.
2C6H18(l) + 25O2 (g)
16CO2(g) + 18H2O(l)
Combustion Reactions – an element or a compound
reacts with oxygen often producing energy in the
form of heat and light..
Many are redox reactions
What Kind of Reaction Is It?
Na2S(aq) + Cd(NO3)2(aq)
CdS(s) + 2NaNO3(aq)
Double Replacement Reaction – involving an
exchange of positive ions between two compounds.
2C6H18(l) + 25O2 (g)
16CO2(g) + 18H2O(l)
Combustion Reactions – an element or a compound
reacts with oxygen often producing energy in the
form of heat and light..
Many are not redox reactions
Identifying Redox Reactions
If the oxidation number of an element in a reacting
species change, then that element has undergone
either oxidation or reduction.
Many reactions in which color changes occur are
redox reactions.
Balancing Redox Equations
Many redox reaction are too complex to be
balanced by trial and error.
Two systematic methods are available to
balance redox reactions
The two methods are based on the fact that the
total number of electrons gained in reduction
must equal the total number of electrons lost in
oxidation.
One method used oxidation number changes,
and the other used half reactions.
Using Oxidation Number Changes to
Balance Redox Equations
Oxidation Number Method
You balance by comparing the increases and the
decreases in oxidation numbers.
+3
-2
+2 -2
Fe2O3(s) + CO(g)
0
+4 -2
Fe(S) + CO2(g)
Step 1 – assign oxidation numbers to all the atoms in
the equation.
Using Oxidation Number Changes to
Balance Redox Equations
Oxidation Number Method
C oxidized (+2)
+3
-2
+2 -2
Fe2O3(s) + CO(g)
0
+4 -2
Fe(S) + CO2(g)
Fe reduced (-3)
Step 2 – Identify which atoms are oxidized and which
are reduced
Using Oxidation Number Changes to
Balance Redox Equations
Oxidation Number Method
(+2) oxidized
+3
-2
+2 -2
Fe2O3(s) + CO(g)
0
+4 -2
Fe(S) + CO2(g)
(-3) reduced
Step 3 – Use one bracketing line to connect the atoms
that undergo oxidation and another such line to
connect those that undergo reduction.
Using Oxidation Number Changes to
Balance Redox Equations
In a balanced redox equation, the total increase in
oxidation number of the species oxidized must be
balance by the total decrease in the oxidation
number of the species reduced.
Using Oxidation Number Changes to
Balance Redox Equations
Oxidation Number Method
3 x (+2) = +6
+3
-2
Fe2O3(s)
+
+2 -2
3CO(g)
0
+4 -2
2Fe(S) + 3CO2(g)
2 x (-3) = -6
Step 4 – Make the total increase in oxidation number
equal to the total decrease in oxidation number by
using appropriate coefficients.
Using Oxidation Number Changes to
Balance Redox Equations
Oxidation Number Method
Fe2O3(s) + 3CO(g)
2Fe(S) + 3CO2(g)
Step 5 – Make sure that the equation is balanced for
both atoms and charge.
Sample Problem
Oxidation Number Method
S oxidized (+4) x 3= +12
+1 +6 -2
+1 -2
0
2K2Cr2O7(aq) + H2O(l) + 3S(s)
+1 -2 +1
+3 -2
KOH(aq) +
+4 -2
2Cr2O3(s) + 3SO2 (g)
Cr reduced (-3) x 4 = -12
4 Cr atoms must be reduced for each 3 S atom that are
oxidized
Sample Problem
Oxidation Number Method
S oxidized (+4) x 3= +12
+1 +6 -2
+1 -2
0
2K2Cr2O7) + 2H2O + 3S
+1 -2 +1
+3 -2
+4 -2
4KOH + 2Cr2O3 + 3SO2
Cr reduced (-3) x 4 = -12
Check the equation and balance by inspection
4 in front of KOH balances potassium
2 in front of H2O balances hydrogen and oxygen.
Sample Problems
Balance each redox equation using the oxidation
number change method
KClO3(s)
2KClO3(s)
KCl (s) + O2 (g)
2KCl (s) + 3O2 (g)
HNO2 (aq) + HI (aq)
NO (g) + I2 (s) + H2O (l)
2HNO2 (aq) + 2HI (aq)
2NO (g) + I2 (s) + 2H2O (l)
Sample Problems
Balance each redox equation using the oxidation
number change method
Bi2S3 + HNO3
Bi2S3 + 8 HNO3
Bi(NO3)3 + NO + S + H2O
2Bi(NO3)3 + 2NO + 3S + 4H2O
SbCl5 + KI
SbCl3 + KCl + I2
SbCl5 + 2KI
SbCl3 + 2KCl + I2
Using Half Reactions to Balance
Redox Equations
Half Reaction Method
Half reaction is an equation showing just the oxidation
or just the reduction that takes place
You write and balance the oxidation and reduction half
reaction separately before combining them into a
balanced redox equation
Then you balance the electrons gained in the reduction
with the electrons lost in the oxidation.
Using Half Reactions to Balance
Redox Equations
Half Reaction Method
S + HNO3
SO2 + NO + H2O (unbalanced)
S + H+ + NO3-
SO2 + NO + H2O
In this case only HNO3 is ionized. The products are
covalent compounds
Step 1 – write the unbalanced equation in ionic form
Using Half Reactions to Balance
Redox Equations
Half Reaction Method
Oxidation Half Reaction
0
+2
S
SO2
Reduction Half Reaction
+5
+2
NO3-
NO
Step 2 – Write separate half reactions for the oxidation
and reduction processes.
Using Half Reactions to Balance
Redox Equations
Half Reaction Method
0
S
+2
SO2
Sulfur is already balanced, but oxygen is not.
The reaction takes place in acid solution, so H2O and
H+ are present and can be used to balance oxygen
and hydrogen as needed.
2H2O + S
SO2 + 4H+
Step 3 – Balance the atoms in the half reactions.
Using Half Reactions to Balance
Redox Equations
Half Reaction Method
+5
+2
NO3-
NO
Nitrogen is already balanced, but oxygen is not.
4H+ + NO3-
NO + 2H2O
Using Half Reactions to Balance
Redox Equations
Half Reaction Method
0
+4
SO2 + 4H+ + 4e -
2H2O + S
S is oxidized going from 0 to +4, a loss of 4 e+5
3e- + 4H+ + NO3-
+2
NO + 2H2O
N is reduced going from +5 to +2, a gain of 3 eStep 4 – Add enough electrons to one side of each half
reaction to balance the charges.
Using Half Reactions to Balance
Redox Equations
Half Reaction Method
0
(x 3)
+4
3SO2 + 12H+ + 12e -
6H2O + 3S
S is oxidized going from 0 to +4, a loss of 4 e+5
(x 4)
12e- + 16H+ + 4NO3-
+2
4NO + 8H2O
N is reduced going from +5 to +2, a gain of 3 eStep 5 – Multiply each half reaction by an appropriate
number to make the numbers of electrons equal in
both.
Using Half Reactions to Balance
Redox Equations
Half Reaction Method
6H2O + 3S
3SO2 + 12H+ + 12e -
12e- + 16H+ + 4NO3-
4NO + 8H2O
6H2O + 3S + 16H+ + 4NO3- + 12e 4NO + 8H2O + 12e –
3S + 4H+ + 4NO3-
3SO2 + 12H+ +
3SO2 + 4NO + 2H2O
Step 6 – Add the balanced half reactions to show an
overall equation and then subtract the terms that
appear in both sides of the equation
Using Half Reactions to Balance
Redox Equations
Half Reaction Method
3S + 4HNO3-
3SO2 + 4NO + 2H2O
Spectator ion – are present during a reaction, but do
not participate in or change during a reaction.
Because none of the ions in the reactants appear in the
products, there are no spectator ions in this
particular example.
Step 7 – Add the spectator ions and balance the
equation.
Sample Problem
Half Reaction Method
KMnO4 + HCl
K+ + MnO4- + H+ + Cl-
MnCl2 + Cl2 + H2O + KCl
Mn2+ + 2Cl- + Cl2 + H2O + K+ +
Cl-
Step 1 – write the unbalanced equation in ionic form
Sample Problem
Half Reaction Method
Reduction Half Reaction
+7
MnO4-
+2
Mn2+
Oxidation Half Reaction
-1
2Cl-
0
Cl2
Step 2 – Write separate half reactions for the oxidation
and reduction processes.
Sample Problem
Reduction Half Reaction
+7
8H+ + MnO4-
+2
Mn2+ + 4H2O
Oxidation Half Reaction
-1
2Cl-
0
Cl2
Solution is acidic, so H2O and H+ ions are used to
balance equation. (If solution is basic, H2O and OHare used)
Step 3 – Balance the atoms in the half reactions.
Sample Problem
Half Reaction Method
+7
+2
5e- + 8H+ + MnO4-
Mn2+ + 4H2O
Mn is reduced going from +7 to +2, a gain of 5 e-1
2Cl-
0
Cl2 + 2e-
Cl is oxidized going from -1 to 0, a loss of 2 eStep 4 – Add enough electrons to one side of each half
reaction to balance the charges.
Sample Problem
Half Reaction Method
+7
+2
(x 2) 10e- + 16H+ + 2MnO4-
2Mn2+ + 8H2O
Mn is reduced going from +7 to +2, a gain of 5 e-1
(x5) 10Cl-
0
5Cl2 + 10e-
Cl is oxidized going from -1 to 0, a loss of 2 eStep 5 – Multiply each half reaction by an appropriate
number to make the numbers of electrons equal in
both.
Sample Problem
Half Reaction Method
10e- + 16H+ + 2MnO42Mn2+ + 8H2O
10Cl5Cl2 + 10e10e- + 16H+ + 2MnO4- + 10Cl5Cl2 + 10e16H+ + 2MnO4- + 10Cl-
2Mn2+ + 8H2O +
2Mn2+ + 8H2O + 5Cl2
Step 6 – Add the balanced half reactions to show an
overall equation and then subtract the terms that
appear in both sides of the equation
Sample Problem
Half Reaction Method
16H+ + 6Cl- + 2MnO4- + 2K+ + 10Cl5Cl2 + 2Mn2+ + 4Cl- + 8H2O + 2K+ + 2Cl16H+ = 6Cl- + 10Cl- (from the HCl in original equation)
2MnO4- = 2K+ (from the KMnO4 in original equation)
2Mn2+ = 4Cl- (from the MnCl2 in original equation)
2K+ + 2Cl- (from the KCl in original equation)
Step 7 – Add the spectator ions and balance the
equation.
Sample Problem
Half Reaction Method
16H+ + 16Cl- + 2MnO4- + 2K+
5Cl2 + 2Mn2+ + 6Cl- + 8H2O + 2K+
Adding spectator and non-spectator Cl- on each side gives
16HCl + 2KMnO4
5Cl2 + 2MnCl2 + 8H2O + 2KCl
Step 7 – Add the spectator ions and balance the
equation.
Sample Problem 2
The following takes place in basic solution. Balance the
equation using the half reaction method
Zn + NO3-
NH3 + Zn(OH4)2-
4Zn + NO3- + 6H2O + 7OH-
4Zn(OH4)2- + NH3
Sample Problem 3
The following takes place in basic solution. Balance the
equation using the half reaction method
Zn + As2O3
6Zn + As2O3 + 9H2O
AsH3 + Zn2+
6Zn2+ + 2AsH3 + 12OH-
Choosing a Balancing Method
In some redox reactions, the same element is both
oxidized and reduced. (called a disproportional
reaction)
3Br2 + 6KOH
5KBr + KBrO3 + 3H2O
Br is reduced from 0 to -1
Br is oxidized from 0 to + 5
Equations like above and equation for reactions that
take place in acidic or alkaline solution are best
balanced by the half reaction method.
End of Chapter 20