Transcript Chapter 7 Lecture

Chapter 7
Chemical Reactions
O2
H2
H2
O
2006, Prentice Hall
react to
form?
+ heat
CHAPTER OUTLINE
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Chemical Reactions
Chemical Equation
Balancing Equations
Types of Chemical Reactions
Activity Series of Metals
Aqueous Reactions
Precipitation Reactions
Neutralization and Other Reactions
Heat in Chemical Reactions
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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.
3
CHEMICAL
REACTIONS
In a chemical reaction, atoms are neither
6 oxygen atoms = 6 oxygen atoms
created, nor destroyed
4
CHEMICAL
REACTIONS
 A chemical reaction can be detected by
one of the following evidences:
1. Change of color
2. Formation of a solid
3. Formation of a gas
4. Exchange of heat with surroundings
5
Evidence of Chemical Change
Emission
of Lightof Heat
Color
Change
Release
or
Absorption
Formation
of Solid
Formation
of aPrecipitate
Gas
Why use Chemical Equations?
1. Shorthand way of describing a reaction
2. Provides information about the reaction
– Formulas of reactants and products
– States of reactants and products
– Relative numbers of reactant and product
molecules that are required
– Can be used to determine amounts of the
reactants and products
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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Symbols Used in Equations
1. energy symbols used above the arrow
for decomposition reactions
–  = heat
– hn = light
– shock = mechanical
– elec = electrical
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
Count and
1C
compare each
element on 4both
H
sides of the
2O
equation
=
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
the
4 Ofor
present
smallest coefficientson
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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TYPES OF
CHEMICAL REACTIONS
 Chemical reactions are can be classified into five
types: Based on what the atoms do
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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Single Displacement
The Zinc Replaces the Copper
Zn(s)  CuCl
2 (aq)
 Cu(s)  ZnCl
2
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)
2 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”.
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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AQUEOUS
REACTIONS
 These
Many substances
ionic solids are
dissolve
calledinelectrolytes.
water to form ions.
NaCl(s)
Na+(aq) + Cl(aq)
H2O
K 2 C rO 4 (s)   
 2 K
H 2O
+
(a q )
+ C rO
24 (a q )
B a (N O 3 ) 2 (s)  2 
 B a (a q ) + 2 N O 3 ( a q )
H O
2+
-
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AQUEOUS
REACTIONS
 When ionic substances dissolve in water they
separate into ions.
K2CrO4
Ba(NO3)2
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AQUEOUS
REACTIONS
 electrolytes are
substances whose
water solution is a
conductor of
electricity
 electrolytes are ions
dissolved in water
(Na+ + Cl-)
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Types of Electrolytes
• salts = water soluble ionic compounds
• acids = form H+1 ions in water solution
– react and dissolve many metals
– strong acid = strong electrolyte, weak acid = weak
electrolyte
• bases = water soluble metal hydroxides
– increases the OH-1 concentration
When will a Salt Dissolve?
• a compound is soluble in a liquid
if it dissolves in that liquid
– NaCl is soluble in water
• a compound is insoluble if a
significant amount does not
dissolve in that liquid
– AgCl is insoluble in water
• though there is a very small amount
dissolved, but not enough to be
significant
AgCl remains solid = precipitate
AQUEOUS
REACTIONS
 Aqueous reactions occur only when one of the
following conditions is present:
1. Formation of a solid:
Precipitation
2. Formation of water:
Neutralization
3. Formation of a gas:
Unstable product
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PRECIPITATION
REACTIONS
 An aqueous chemical reaction that produces a
solid as one of its products is called a
precipitation reaction.
 The insoluble solid formed in these reactions is
called a precipitate.
K 2 C rO 4 (aq ) + B a(N O 3 ) 2 (aq )  
 B aC rO 4 (s) + 2 K N O 3 (aq )
Precipitate
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Example of a Precipitation Reaction
Pb(NO3)2(aq) + 2 KI(aq)  2 KNO3(aq) + PbI2(s)
Let’s Look Closer at PbI2 Formation
Pb(NO3)2(aq) + 2 KI(aq)  2 KNO3(aq) + PbI2(s)
SOLUBILITY
RULES
 Chemists use a set of solubility rules to predict
No exceptions
S whetherNO
3
a product is soluble or insoluble.
O
Na+, K+
No exceptions
L
+
NH4
U
+
Except
those
containing
Ag
B
Cl, Br, I
, Pb2+
L
Except those containing
2
E
SO4
Ba2+ , Pb2+, Ca2+
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SOLUBILITY
RULES
I
N
S
O
L
S2, CO32
PO43
Except those containing
Na+ , K+, NH4+
OH
Except those containing
Na+ , K+, Ca2+, NH4+
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Example 1:
Write balanced equations for each reactions shown
below. Indicate if no reaction occurs.
NaCl (aq) + AgNO3 (aq)  NaNO3 (?) + AgCl (?)
NaCl (aq) + AgNO3 (aq)  NaNO3 (aq) + AgCl (s)
soluble
precipitate
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Example 2:
Write balanced equations for each reactions shown
below. Indicate if no reaction occurs.
NH4Cl (aq) + KNO3 (aq)  NH4NO3 (?) + KCl (?)
No4NO
Reaction
NH4Cl (aq) + KNO3 (aq)  NH
3 (aq) + KCl (aq)
soluble
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Example 3:
Write balanced equations for each reactions shown
below. Indicate if no reaction occurs.
PbCl2 (aq) + 2 NaI (aq)  PbI2 (?) + 2 NaCl (?)
PbCl2 (aq) + 2 NaI (aq)  PbI2 (s) + 2 NaCl (aq)
precipitate
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Molecular, Complete Ionic, and Net
Ionic Equations
A molecular equation is a chemical equation
showing the complete, neutral formulas for
every compound in a reaction.
A complete ionic equation is a chemical
equation showing all of the species as they
are actually present in solution.
A net ionic equation is an equation showing
only the species that actually participate in
the reaction.
© 2012 Pearson Education, Inc.
Writing Chemical Equations for Reactions in Solution:
Molecular and Complete Ionic Equations
• A molecular equation is an equation showing the
complete neutral formulas for every compound in the
reaction.
• Complete ionic equations show aqueous ionic
compounds that normally dissociate in solution as they
are actually present in solution.
• When writing complete ionic equations, separate only
aqueous ionic compounds into their constituent ions.
• Do NOT separate solid, liquid, or gaseous
compounds.
© 2012 Pearson Education, Inc.
Writing Chemical Equations for Reactions in Solution:
Net Ionic Equations
• In the complete ionic equation, some of the ions
in solution appear unchanged on both sides of
the equation.
• These ions are called spectator ions because
they do not participate in the reaction.
© 2012 Pearson Education, Inc.
Writing Chemical Equations for Reactions in Solution:
Proper Net Ionic Equations
• To simplify the equation, and to more clearly
show what is happening, spectator ions can
be omitted.
• Equations such as this one, which show only
the species that actually participate in the
reaction, are called net ionic equations.
Ag+(aq) + Cl− (aq)  AgCl(s)
© 2012 Pearson Education, Inc.
NEUTRALIZATION
REACTIONS
 The
Saltsmost
are ionic
important
substances
reaction
withofthe
acids
cation
and
bases is called
donated
from the
neutralization.
base and the anion donated
thereactions
acid.
 from
In these
an acid combines with a
base to form a salt and water.
H C l (aq ) + N aO H (aq ) ¾ ¾® N aC l (aq ) + H 2 O (l)
Acid
Base
Salt
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Examples:
Write balanced equations for each of the neutralization reactions shown below:
2 HNO3 + Ba(OH)2  Ba(NO3)2 + 2 H2O
H2SO4 + 2 NaOH 
Na2SO4 + 2 H2O
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GAS FORMING
REACTIONS
 Some chemical reactions produce gas because
one of the products formed in the reaction is
unstable.
 Three such products are:
Carbonic acid: H2CO3 (aq)  CO2 (g) + H2O (l)
Sulfurous acid: H2SO3 (aq)  SO2 (g) + H2O (l)
Ammonium:
NH4OH (aq)  NH3 (g) + H2O (l)
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GAS FORMING
REACTIONS
 When either of these products appears in a
chemical reaction, they should be replaced with
their decomposition products.
22HNO
HCl +
Na
K22CO
SO3  2 KNO
NaCl 3++ HH2CO
3+
2SO33
2 HNO
2 HCl3 + Na2SO
CO33 
 22 KNO
NaCl3 ++ CO
SO2 (g)
(g)++ H
H22O
O (l)
(l)
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Enthalpy: A Measure of the Heat
Evolved or Absorbed in a Reaction
• Chemical reactions can be exothermic (they
emit thermal energy when they occur).
• Chemical reactions can be endothermic
(they absorb thermal energy when they
occur).
• The amount of thermal energy emitted or
absorbed by a chemical reaction, under
conditions of constant pressure (which are
common for most everyday reactions), can
be quantified with a function called enthalpy.
© 2012 Pearson Education, Inc.
Enthalpy: A Measure of the Heat
Evolved or Absorbed in a Reaction
• We define the enthalpy of reaction,
ΔHrxn, as the amount of thermal energy (or
heat) that flows when a reaction occurs at
constant pressure.
© 2012 Pearson Education, Inc.
Sign of ΔHrxn
• The sign of ΔHrxn (positive or negative)
depends on the direction in which thermal
energy flows when the reaction occurs.
• Energy flowing out of the chemical system is
like a withdrawal and carries a negative sign.
• Energy flowing into the system is like a
deposit and carries a positive sign.
© 2012 Pearson Education, Inc.
Exothermic and Endothermic
reactions
• (a) In an exothermic reaction, energy is released
into the surroundings. (b) In an endothermic
reaction, energy is absorbed from the surroundings.
© 2012 Pearson Education, Inc.
HEAT IN CHEMICAL
REACTIONS
 Reactions
In exothermic
that reaction,
release heat
heatare
is produced
classified as
and can
exothermic.
be
written as a product.
 In
Reactions
endothermic
that absorb
reaction,
heat
heat
areisclassified
requiredas
and
can
be written as a reactant.
endothermic.
Endothermic
H2 (g) + Cl2 (g)  2 HCl (g) + 185 kJ
N2 (g) +Exothermic
O2 (g) + 181 kJ  2 NO (g)
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Sign of ΔHrxn
• When thermal energy flows out of the reaction and
into the surroundings it is a ??? reaction and has a
+ or – enthalpy?
• The enthalpy of reaction for the combustion of
CH4, the main component in natural gas:
• The magnitude of ΔHrxn tells us that 802.3 kJ of
heat are emitted when 1 mol CH4 reacts with 2
mol O2.
© 2012 Pearson Education, Inc.
Stoichiometry of ΔHrxn
• The amount of heat emitted or absorbed
when a chemical reaction occurs depends
on the amounts of reactants that actually
react.
• We usually specify ΔHrxn in combination
with the balanced chemical equation for
the reaction.
• The magnitude of ΔHrxn is for the
stoichiometric amounts of reactants and
products for the reaction as written.
© 2012 Pearson Education, Inc.
Stoichiometry of ΔHrxn
• The balanced equation and ΔHrxn for the
combustion of propane is:
• When 1 mole of C3H8 reacts with 5 moles of O2 to
form 3 moles of CO2 and 4 moles of H2O, 2044 kJ
of heat are emitted.
© 2012 Pearson Education, Inc.
Example 1:
• An gas tank in a home barbecue contains
11.8 x 103 g of propane (C3H8).
• Calculate the heat (in kJ) associated with
the complete combustion of all of the
propane in the tank.
© 2012 Pearson Education, Inc.
Example 1:
© 2012 Pearson Education, Inc.
Classifying Reactions
• Also we can classify reactions by what
happens:
• Redox reactions are the exchange of e• Redox are all reactions except?
Note
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
of electrons.
a nail or the reaction within your car
batteries are two examples of redox reactions.
64
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
Reduction Is Gain of
electrons
OIL
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 lose
of electrons
calcium to
form cations,
sulfide
from calcium
and areand
therefore
sulfur oxidized,
while atomsCa
of +non-metals
gain electrons to
S
CaS
form anions, and are therefore reduced.
Ca2+ + 2
e
S + 2 eS2-
Ca
Oxidation
Reduction
 Therefore, the formation of calcium sulfide
involves two half-reactions that occur
simultaneously, one an oxidation and the other
66
a reduction.
COMBUSTION
 A reaction that involves oxygen as a reactant
and produces large amounts
of heat is classified
occurs in
as a combustion reaction. the cylinders
of the engine
 Combustion reactions are a subclass of
Oxidation-Reduction reactions
2 C8H18(g) + 25 O2(g)  16 CO2(g) + 18 H2O(g)
67
Combustion Products
• predicting the products of a combustion
reaction; simply combine each element in
the other reactant with oxygen
Reactant
Combustion
Product
contains C
CO2(g)
contains H
H2O(g)
contains S
SO2(g)
contains N
NO(g) or NO2(g)
contains metal
M2On(s)
Combustion Reactions
• combustion reactions are always exothermic
• in combustion reactions, O2 combines with the
elements in another reactant to make the products
4 Fe(s) + 3 O2(g) → 2 Fe2O3(s) + energy
CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g) + energy
The flame on a gas stove results from the oxidation of carbon in natural gas.
Reverse of Combustion Reactions
• since combustion reactions are exothermic,
their reverse reactions are endothermic
• the reverse of a combustion reaction involves
the production of O2
energy + 2 Fe2O3(s) → 4 Fe(s) + 3 O2(g)
energy + CO2(g) + 2 H2O(g) → CH4(g) + 2 O2(g)
• reactions in which O2 is gained or lost are
redox reactions
REDOX 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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REDOX IN
BIOLOGICAL SYSTEMS
 The formaldehyde is further oxidized to formic
acid and finally carbon dioxide and water by
the
Oxidation
following reactions:
(gain of
2 H2CO + O2
formaldehyde
2 H2CO2 + O2
2H2CO2
oxygen)
formic acid
CO2 + H2O
formic acid
72
REDOX IN
BIOLOGICAL SYSTEMS
 In
Themany
oxidation
biochemical
of a typical
oxidation-reduction
biochemical molecule
reactions,
can
the transfer
involve
the transfer
of hydrogen
of two
atoms
hydrogen
produces
atoms
energy
to a in
the cells.
proton
acceptor such as coenzyme FAD to produce
its reduced form FADH2.
73
REDOX IN
BIOLOGICAL SYSTEMS
 In summary, the particular definition of oxidationreduction depends on the process that occurs in the
reaction.
74
THE END
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