Chapter 6A
Chemical
Reactions
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CHAPTER OUTLINE
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Chemical Reactions
Chemical Equation
Balancing Equations
Types of Chemical Reactions
Double Replacement Reactions
Oxidation-Reduction Reactions
Redox in Biological Systems
Activity Series of Metals
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CHEMICAL
REACTIONS
 A chemical reaction is a rearrangement of
atoms in which some of the original bonds are
broken and new bonds are formed to give
different chemical structures.
 In a chemical reaction, atoms are neither
created, nor destroyed.
 A chemical reaction, as described above, is
supported by Dalton’s postulates.
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CHEMICAL
REACTIONS
In a chemical reaction, atoms are neither
6 oxygen atoms = 6 oxygen atoms
created, nor destroyed
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CHEMICAL
REACTIONS
 A chemical reaction can be detected by one of
the following evidences:
1. Change of color (formation of a solid)
2. Formation of a gas
3. Exchange of heat with surroundings
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CHEMICAL
EQUATIONS
 A chemical equation is a shorthand expression
for a chemical reaction.
Word equation:
Aluminum combines with ferric oxide to form
iron and aluminum oxide.
Chemical equation:
Al + Fe2O3  Fe + Al2O3
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CHEMICAL
EQUATIONS
 Reactants are separated from products by an
arrow.
Al + Fe2O3  Fe + Al2O3
 Coefficients are placed in front of substances
to balance the equation.
2 Al + Fe2O3  2 Fe + Al2O3
Subscripts
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CHEMICAL
EQUATIONS
 Reaction conditions are placed over the arrow.

Al + Fe2O3  Fe + Al2O3
heat
 The physical state of the substances are
indicated by the symbols (s), (l), (g), (aq).

2 Al (s) + Fe2O3 (s)  2 Fe (l) + Al2O3 (s)
solid
liquid
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BALANCING
EQUATIONS
 A balanced equation contains the same number
of atoms on each side of the equation, and
therefore obeys the law of conservation of
mass.
 Many equations are balanced by trial and
error; but it must be remembered that
coefficients can be changed in order to balance
an equation, but not subscripts of a correct
formula.
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BALANCING
EQUATIONS
 The general procedure for balancing equations is:
Write the unbalanced equation:
CH4 + O2  CO2 + H2O
Make sure the
formula for each
substance is correct
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BALANCING
EQUATIONS
 The general procedure for balancing equations is:
Balance by inspection:
CH4 + O2  CO2 + H2O
1C
Count and compare
each element on
4H
both sides of the
equation2 O
=
1C

2H

3O
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BALANCING
EQUATIONS
 Balance elements that appear only in one
substance first.
Balance H
4 H present on
each side
CH4 + O2  CO2 + H2O
1 CH4 + O2  CO2 + 2 H2O
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BALANCING
EQUATIONS
Balance O
When finally done, check
for the on
4 O present
smallest coefficients possible
each side
1 CH4 + O2  CO2 + 2 H2O
1 CH4 + 2 O2  CO2 + 2 H2O
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Examples:
2 AgNO
AgNO33 ++ H
H22SS 
 Ag
Ag22SS ++ 2HNO
HNO
3 3
2 Al(OH)
Al(OH)3 3++3 H2SO4  Al2(SO4)3 + 6H2H
O2O
FeFe
4H
H22 
 3Fe
Fe ++ H42H
O2O
3O
3O
4 4+ +
2 C4CH410
H10+ +
13 O
O22 
 8CO
CO
10
2 2+ + H
2OH2O
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CONCEPT
CHECK
 If red spheres represent oxygen atoms and blue
spheres represent nitrogen atoms, write a
balanced equation for the reaction shown below.
2 NO + O2  2 NO2
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TYPES OF
CHEMICAL REACTIONS
 Chemical reactions are classified into five types:
1. Synthesis or combination
2. Decomposition
3. Single replacement
4. Double replacement
5. Combustion
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SYNTHESIS or
COMBINATION
 In these reactions, 2 elements or compounds
combine to form another compound.
A + B  AB
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DECOMPOSITION
 In these reactions, a compound breaks up to form
2 elements or simpler compound.
AB  A +B
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SINGLE
REPLACEMENT
 In these reactions, a more reactive element
replaces a less reactive element in a compound.
A + BC  B + AC
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DOUBLE
REPLACEMENT

 The
In these
cation
reactions,
from onetwo
compound
compounds
replaces
combine
the cation
to
in
form
another
two new
compound.
compounds.
+
+
AB + CD  AD + CB
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COMBUSTION
 A reaction that involves oxygen as a reactant
and produces large amounts of heat is classified
as a combustion reaction.
CH4 (g) + 2 O2 (g)  CO2 (g) + 2 H2O (g)
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Examples:
Classify each of the reactions below:
1.
2.
3.
4.
Decomposition
Mg + CuCl2  MgCl2 + Cu
CaCO3  CaO
+ CO2
Synthesis
Single replacement
2 HCl + Ca(OH)2  CaCl2 + 2 H2O
reactive
4 FeMg
+ 3isOmore
Fe2O3 than Cu
2 2
Double replacement
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DOUBLE REPLACEMENT
REACTIONS
 Double replacement reactions can be subdivided
into one the following subgroups:
1. Precipitation:
formation of a solid
2. Neutralization: formation of water
3. Unstable product: formation of a gas
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PRECIPITATION
REACTIONS
 In these reactions one of the products formed is an
insoluble solid called a precipitate.
 For example, when solutions of potassium chromate,
K2CrO4 , and barium nitrate, Ba(NO3)2 , are
combined an insoluble salt barium chromate,
BaCrO4 , is formed.
K2CrO4 (aq) + Ba(NO3)2 (aq)
BaCrO4 (s) + 2 KNO3 (aq)
precipitate
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NEUTRALIZATION
REACTIONS
 The
Saltsmost
are ionic
important
substances
reaction
withofthe
acids
cation
anddonated
bases is
calledthe
from
neutralization.
base and the anion donated from the acid.
 In the
these
laboratory,
reactions an
neutralization
acid combines
reactions
with a are
base to
form a salt
observed
byand
an increase
water. in temperature
(exothermic reaction).
HCl (aq) + NaOH (aq) ¾ ¾
® NaCl (aq) + H 2O (l)
Acid
Base
Salt
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UNSTABLE PRODUCT
REACTIONS
 Some chemical reactions produce gas because one
of the products formed in the reaction is unstable.
 Two such products are:
Carbonic acid: H2CO3 (aq)  CO2 (g) + H2O (l)
Sulfurous acid: H2SO3 (aq)  SO2 (g) + H2O (l)
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UNSTABLE PRODUCT
REACTIONS
 When either of these products appears in a chemical
reaction, they should be replaced with their
decomposition products.
2 HCl + Na2CO3  2 NaCl + H2CO3
2 HCl + Na2CO3  2 NaCl + CO2 (g) + H2O (l)
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Examples:
Complete and balance each neutralization reaction
below:
2 HNO3 + Ba(OH)2 
Ba(NO3)2 + 2 H2O
H2SO4 + 2 NaOH 
Na2SO4 + 2 H2O
HC2H3O2 + KOH  KC2H3O2 + H2O
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OXIDATION-REDUCTION
REACTIONS



In an oxidation-reduction
Reactions
known as oxidation
reaction,
and reduction
electrons are
(redox) havefrom
transferred
many
one
important
substance
applications
to another.in our
everyday
If
one substance
lives. loses electrons, another substance
Rusting
must
gain
ofelectrons.
a nail or the reaction within your car
batteries are two examples of redox reactions.
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OXIDATION-REDUCTION
REACTIONS



Oxidation is defined as loss of electrons, and
reduction is defined as gain of electrons.
One way to remember these definitions is to use
the following mnemonic:
Oxidation Is Loss of electrons
OIL
Reduction Is Gain of electrons
RIG
Combination, decomposition, single replacement
and combustion reactions are all examples of
redox reactions.
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OXIDATION-REDUCTION
REACTIONS
 In
Forgeneral,
example,
atoms
in the
offormation
metals loseofelectrons
calcium sulfide
to form
cations,
from
calcium
and are
and
therefore
sulfur oxidized, while atoms of
non-metals gain
toCaS
form anions, and are
Caelectrons
+S
therefore reduced.
Ca
Ca2+ + 2 e-
S + 2 e-
S2-
Oxidation
Reduction
 Therefore, the formation of calcium sulfide involves
two half-reactions that occur simultaneously, one an
oxidation and the other a reduction.
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OXIDATION-REDUCTION
REACTIONS
 Similarly, in the reaction of magnesium metal with
hydrochloric acid
Mg + 2 HCl
Mg
MgCl2 + H2
Mg2+ + 2 e-
2 H+ + 2 e-
H2
Oxidation
Reduction
In every oxidation-reduction reaction, the number of electrons
lost must be equal to the number of electrons gained.
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OXIDATION-REDUCTION IN
BIOLOGICAL SYSTEMS
 Many important biological reactions involve
oxidation and reduction.
 In these reactions, oxidation involves addition of
oxygen or loss of hydrogen,
and reduction involves
Oxidation
loss of oxygen or gain of hydrogen.
(loss of
 For example, poisonous methyl
alcohol is
hydrogen)
metabolized by the body by the following reaction:
CH3OH
methyl alcohol
H2CO + 2H•
formaldehyde
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OXIDATION-REDUCTION IN
BIOLOGICAL SYSTEMS
 The formaldehyde is further oxidized to formic
acid and finally carbon dioxide and water by the
Oxidation
following reactions:
(gain of
oxygen)
2 H2CO + O2
2H2CO2
formaldehyde
2 H2CO2 + O2
formic acid
CO2 + H2O
formic acid
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OXIDATION-REDUCTION IN
BIOLOGICAL SYSTEMS
 In many biochemical oxidation-reduction reactions,
the transfer of hydrogen atoms produces energy in
the cells.
 For example, cellular respiration is an oxidationreduction process that transfers energy from the
bonds in glucose to form ATP.
C6H12O6 + 6 O2
6 CO2 + 6 H2O + ATP + Heat
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OXIDATION-REDUCTION IN
BIOLOGICAL SYSTEMS
Loss of hydrogen atoms
(becomes oxidized)
C6H12O6
6 O2
6 CO2
6 H2O
Glucose
Gain of hydrogen atoms
(becomes reduced)
ATP
 Heat
OXIDATION-REDUCTION IN
BIOLOGICAL SYSTEMS
 The oxidation of a typical biochemical molecule can
involve the transfer of hydrogen atoms to a proton
acceptor such as coenzyme FAD to produce its
reduced form FADH2.
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REDOX IN
BIOLOGICAL SYSTEMS
 In summary, the particular definition of oxidationreduction depends on the process that occurs in the
reaction.
 A summary of these definitions appears below:
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Example 1:
Linoleic acid, an unsaturated fatty acid, can be converted
to a saturated fatty acid by the reaction shown below. Is
linoleic acid oxidized or reduced in this reaction?
C18H32 O2 + 2 H2
C18H36 O2
Gain of
Reduction
hydrogen
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Example 2:
The reaction of succinic acid (C4H6O4) provides energy
for the ATP synthesis and is shown below:
C4H6 O4
FAD + 2 H•
C4H4 O4 + 2 H•
FADH2
a) Is succinic acid oxidized or reduced? oxidation
Loss of
b) Is FAD oxidized or reduced? reduction
hydrogen
Gain of
hydrogen
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ENZYMES IN
BIOLOGICAL SYSTEMS
 In biochemical reactions, enzymes are necessary to
oxidize glucose and other foods.
 For example, oxidation
of glucose involves the
transfer of hydrogen
atoms and electrons to
an enzyme, such as
NAD+ to produce its
reduced form NADH.
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ENZYMES IN
BIOLOGICAL SYSTEMS
 Similarly, oxidation of methanol involves transfer
of 2 hydrogen atoms and 2 electrons to NAD+ to
form the reduced form NADH.
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ENZYMES IN
BIOLOGICAL SYSTEMS
 Molecules such as NAD+ are called “electron
carriers” since they carry electrons in their reduced
form.
 The electron carriers collectively are called electron
transport chain.
 As electrons are transported down the chain, ATP
is generated.
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ENZYMES IN
BIOLOGICAL SYSTEMS
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ACTIVITY SERIES
OF METALS
 Activity series is a listing of metallic elements in
descending order of reactivity.
 Hydrogen is also included in the series since it
behaves similar to metals.
 Activity series tables are available in textbooks
and other sources.
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ACTIVITY SERIES
OF METALS
 Elements listed higher will
displace any elements listed below
them.
 For example Na will displace any
elements listed below it from one
of its compounds.
2 Na (s) + MgCl2 (aq)  2 NaCl (aq) + Mg (s)
Na (s) + AgCl (aq)  NaCl (aq) + Ag (s)
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ACTIVITY SERIES
OF METALS
 Elements listed lower will not
displace any elements listed above
them.
 For example Ag cannot displace
any elements listed above it from
one of its compounds.
Ag (s) + CuCl2 (aq)  No Reaction
Ag (s) + HCl (aq)  No Reaction
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Example 1:
Use activity series to complete each reaction below.
If no reaction occurs, write “No Reaction”.
Pb (s) + 2 HCl (aq)  PbCl2 (aq) + H2 (g)Metals
Pb is more
reactive
than H
Fe
Ni
Sn
Pb
H
Cu
Ag
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Example 2:
Use activity series to complete each reaction below.
If no reaction occurs, write “No Reaction”.
Zn (s) + MgCl2 (aq)  No Reaction
Zn is less
reactive
than Mg
Metals
Na
Mg
Al
Zn
Fe
Ni
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Example 3:
Use activity series to complete each reaction below.
If no reaction occurs, write “No Reaction”.
Ni (s) + CuCl2 (aq)  NiCl2 (aq) + Cu (s)Metals
Ni is more
reactive
than Cu
Fe
Ni
Sn
Pb
H
Cu
Ag
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Example 4:
Use activity series to complete each reaction below.
If no reaction occurs, write “No Reaction”.
3 Mg (s) + 2 AlCl3 (aq)  3 MgCl2 (aq) + 2Metals
Al (s)
Mg is more
reactive
than Al
Na
Mg
Al
Zn
Fe
Ni
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THE END
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