Download Chapter 6A Chemical Reactions CHAPTER OUTLINE

4/19/15
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
q  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.
q  In a chemical reaction, atoms are neither
created, nor destroyed.
q  A chemical reaction, as described above, is
supported by Dalton’s postulates.
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CHEMICAL
REACTIONS
a chemical reaction, atoms are neither
6In
oxygen
atoms = 6 oxygen atoms
created, nor destroyed
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CHEMICAL
REACTIONS
q  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
q  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
q  Reactants are separated from products by an
arrow.
Al + Fe2O3 → Fe + Al2O3
q  Coefficients are placed in front of substances
to balance the equation.
2 Al + Fe2O3 → 2 Fe + Al2O3
Subscripts
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CHEMICAL
EQUATIONS
q  Reaction conditions are placed over the arrow.
Δ
Al + Fe2O3 →
Fe + Al2O3
heat
q  The physical state of the substances are
indicated by the symbols (s), (l), (g), (aq).
Δ
2 Al (s) + Fe2O3 (s) →
2 Fe (l) + Al2O3 (s)
solid
liquid
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BALANCING
EQUATIONS
q  A balanced equation contains the same number
of atoms on each side of the equation, and
therefore obeys the law of conservation of
mass.
q  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
q  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
q  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
q  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
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2 2+ + H
2OH2O
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CONCEPT
CHECK
q  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
q  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
q  In these reactions, 2 elements or compounds
combine to form another compound.
A + B → AB
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DECOMPOSITION
q  In these reactions, a compound breaks up to form
2 elements or simpler compound.
AB → A +B
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SINGLE
REPLACEMENT
q  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
q 
q  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
q  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
q  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
q  In these reactions one of the products formed is an
insoluble solid called a precipitate.
q  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
q  The
Saltsmost
are ionic
important
substances
reaction
withofthe
acids
cation
anddonated
bases is
calledthe
from
neutralization.
base and the anion donated from the acid.
q  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) æ æÆ NaCl (aq) + H 2O (l)
Acid
Base
Salt
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UNSTABLE PRODUCT
REACTIONS
q  Some chemical reactions produce gas because one
of the products formed in the reaction is unstable.
q  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
q  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
q  Reactions
In an oxidation-reduction
known as oxidation
reaction,
and reduction
electrons are
(redox) havefrom
transferred
many
one
important
substance
applications
to another.in our
q  everyday
If one substance
lives. loses electrons, another substance
q  must
Rusting
gain
ofelectrons.
a nail or the reaction within your car
batteries are two examples of redox reactions.
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OXIDATION-REDUCTION
REACTIONS
q  Oxidation is defined as loss of electrons, and
reduction is defined as gain of electrons.
q  One way to remember these definitions is to use
the following mnemonic:
Oxidation Is Loss of electrons
OIL
Reduction Is Gain of electrons
RIG
q  Combination, decomposition, single replacement
and combustion reactions are all examples of
redox reactions.
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OXIDATION-REDUCTION
REACTIONS
q  In
Forgeneral,
example,
atoms
in the
offormation
metals loseofelectrons
calcium to
sulfide
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
q  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
q  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
q  Many important biological reactions involve
oxidation and reduction.
q  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
q  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
q  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
q  In many biochemical oxidation-reduction reactions,
the transfer of hydrogen atoms produces energy in
the cells.
q  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
Glucose
6 O2
6 CO2
6 H 2O
Gain of hydrogen atoms
(becomes reduced)
ATP
+ Heat
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OXIDATION-REDUCTION IN
BIOLOGICAL SYSTEMS
q  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
q  In summary, the particular definition of oxidationreduction depends on the process that occurs in the
reaction.
q  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
q  In biochemical reactions, enzymes are necessary to
oxidize glucose and other foods.
q  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
q  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
q  Molecules such as NAD+ are called “electron
carriers” since they carry electrons in their reduced
form.
q  The electron carriers collectively are called electron
transport chain.
q  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
q  Activity series is a listing of metallic elements in
descending order of reactivity.
q  Hydrogen is also included in the series since it
behaves similar to metals.
q  Activity series tables are available in textbooks
and other sources.
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ACTIVITY SERIES
OF METALS
q  Elements listed higher will
displace any elements listed below
them.
q  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
q  Elements listed lower will not
displace any elements listed above
them.
q  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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