OBJECTIVES:

❑ Describe chemical
reaction.
❑ Differentiate the types of
chemical reaction.
❑ Explain chemical equation.
❑ Balance chemical equation.
1
Chemical Reactions

O2
react to
form?
H2

H2O + heat
 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
Release Color
or of Lightof Heat
Change
Absorption
Formation of Solid
Formation of aPrecipitate
Gas
 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
7
 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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 DOUBLE
REPLACEMENT
❑ Precipitation reaction yields precipitate (solid
substance that settles at the bottom) when 2
solutions are mixed.

❑ Neutralization is a reaction that occurs between an

acid and a base that forms salt and water.

12
 Examples:
Classify each of the reactions below:
Decomposition
1. Mg + CuCl2 → MgCl2 + Cu
2. CaCO3 → CaO + CO2
Synthesis
Single replacement
3. 2 HCl + Ca(OH)2 → CaCl2 + 2 H2O
4. 4 Fe +Mg
3 Ois2 more
→ 2 Fe O
reactive
2 3 than Cu
Double replacement

13
 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
14
 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.

Δ
Al + Fe2O3 → Fe + Al2O3
heat
❑ 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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 Symbols Used in Equations
1. energy symbols used above the arrow
for decomposition reactions
– Δ = heat
– hν = 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
19
 BALANCING
EQUATIONS
❑ The general procedure for balancing equations is:

Balance by inspection:

CH4 + O2 → CO2 + H2O

Count and
1C = 1C
compare each
element on 4both
H  2H
sides of the
equation
2O  3O
20
 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

21
 BALANCING
EQUATIONS

Balance O

When finally done, check

4 Ofor the
present
smallest coefficientson
possible
each side

1 CH4 + O2 → CO2 + 2 H2O

1 CH4 + 2 O2 → CO2 + 2 H2O

22
Examples:

2 AgNO
AgNO33 ++ H
H22SS →
→ Ag
Ag22SS ++ 2HNO
HNO3 3

2 Al(OH)
Al(OH)3 3++3 H2SO4 → Al2(SO4)3 + 6H2H
O2O

FeFeOO + +4 H
3 34 4
H22 →
→ 3Fe
Fe ++ H42H
O2O

2 C4CH410
H10+ +
13 O
O22 →
→ 8CO
CO + + H10
2 2 2
OH2O

23
 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)

24
Single Displacement
The Zinc Replaces the Copper
 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.

26
 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)

27
 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

28
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

Fe
Ni
Sn
Pb is more Pb
reactive H
than H Cu
Ag
29
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

Fe
Ni
Sn
Ni is more Pb
reactive H
than Cu Cu
Ag
30
 AQUEOUS
REACTIONS
❑ Many substances
These ionic solids are
dissolve
calledinelectrolytes.
water to form ions.

NaCl(s) H2O
Na+(aq) + Cl(aq)

31
 AQUEOUS
REACTIONS
❑ When ionic substances dissolve in water they
separate into ions.

K2CrO4 Ba(NO3)2

32
 AQUEOUS
REACTIONS

□ electrolytes are
substances whose
water solution is a
conductor of
electricity
□ electrolytes are ions
dissolved in water
(Na+ + Cl-)

33
 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

36
 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.

Precipitate
37
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

S whetherNO 3
No exceptions
a product is soluble or insoluble.
O
Na+, K+
L + No exceptions
NH4
U
+
Except those containing Ag
B Cl−, Br−, I−
, Pb2+
L
E 2−
Except those containing
SO4 Ba2+ , Pb2+, Ca2+
40
 SOLUBILITY
RULES

I S2−, CO32− Except those containing

N PO43− Na+ , K+, NH4+
S
O − Except those containing
OH
L Na+ , K+, Ca2+, NH4+

41
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
42
 Example 2:
Write balanced equations for each reactions shown
below. Indicate if no reaction occurs.

NH4Cl (aq) + KNO3 (aq) → NH4NO3 (?) + KCl (?)

NH4Cl (aq) + KNO3 (aq) → NH

No4NO
Reaction
3
(aq) + KCl (aq)

soluble

43
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

44
 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
❑ Saltsmost
The are ionic
important
substances
reaction
withofthe
acids
cation
and
bases is called
donated from the
neutralization.
base and the anion donated
❑ from thereactions
In these acid. an acid combines with a
base to form a salt and water.

Acid Base Salt

49
Examples:
Write balanced equations for each of the
neutral-ization reactions shown below:

2 HNO3 + Ba(OH)2 → Ba(NO3)2 + 2 H2O

H2SO4 + 2 NaOH → Na2SO4 + 2 H2O

50
 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)

51
 GAS FORMING
REACTIONS
❑ When either of these products appears in a
chemical reaction, they should be replaced with
their decomposition products.

22HNO
HCl +
3
+Na
K22CO
SO3 → 2 KNO
NaCl 3++ HH2CO
2
SO33

2 HNO
2 HCl3 + Na2SO
CO33 →
→ 22 KNO
NaCl3 ++ CO
SO2 (g)
(g)++ H
H22O
O (l)
(l)

52
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
❑ In exothermic
Reactions that reaction,
release heat
heatare
is produced
classified as
and can
exothermic.
be written as a product.
❑ Reactions
In endothermic
that absorb
reaction,
heat
heat
areisclassified
requiredas and can
endothermic.
be written as a reactant.

Endothermic
H2 (g) + Cl2 (g) → 2 HCl (g) + 185 kJ

N2 (g) +Exothermic
O2 (g) + 181 kJ → 2 NO (g)

57
 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
O 2.

© 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 oxidationreaction,
and reduction
electrons are
(redox) havefrom
transferred many one
important
substanceapplications
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.

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 OIL
electrons
Reduction Is Gain of RIG
❑ electrons
Combination, decomposition, single
replacement and combustion reactions are
all examples of redox reactions.
65
 OXIDATION-REDUCTION
REACTIONS
❑ Forgeneral,
In 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
S gain electrons to
CaS
form anions, and are therefore reduced.
Ca Ca2+ + 2 e- Oxidation
S + 2 e- S2- Reduction
❑ Therefore, the formation of calcium sulfide
involves two half-reactions that occur
simultaneously, one an oxidation and the other
a reduction. 66
 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 H2CO + 2H•

methyl alcohol formaldehyde

71
 REDOX IN
BIOLOGICAL SYSTEMS
❑ The formaldehyde is further oxidized to formic
acid and finally carbon dioxide and water byOxidation
the
following reactions: (gain of
oxygen)
2 H2CO + O2 2H2CO2
formaldehyde formic acid

2 H2CO2 + O2 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
oxidation-reduction depends on the process that
occurs in the reaction.

74
THE END

75