Chemistry

1. Chapter 7
Chemical Reactions
2006, Prentice Hall
H2
O2
H2
O
react to
form?
+ heat

2. 2
CHAPTER OUTLINE
 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

3. 3
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.

4. 4
CHEMICAL
REACTIONS
6 oxygen atoms 6 oxygen atoms
=
In a chemical reaction, atoms are neither
created, nor destroyed

5. 5
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

6. Evidence of Chemical Change
Color Change
Formation of Solid Precipitate
Formation of a Gas
Emission of Light
Release or Absorption of Heat

7. 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

8. 8
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

9. 9
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

10. 10
CHEMICAL
EQUATIONS
 Reaction conditions are placed over the arrow.
Al + Fe2O3  Fe + Al2O3
 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)

heat
solid liquid


11. Symbols Used in Equations
1. energy symbols used above the arrow
for decomposition reactions
–  = heat
– hn = light
– shock = mechanical
– elec = electrical

12. 12
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.

13. 13
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

14. 14
BALANCING
EQUATIONS
 The general procedure for balancing equations is:
Balance by inspection:
CH4 + O2  CO2 + H2O
Count and
compare each
element on both
sides of the
equation
1 C = 1 C
4 H  2 H
2 O  3 O

15. 15
BALANCING
EQUATIONS
 Balance elements that appear only in one
substance first.
Balance H
CH4 + O2  CO2 + H2O
1 CH4 + O2  CO2 + 2 H2O
4 H present
on each side

16. 16
BALANCING
EQUATIONS
Balance O
1 CH4 + O2  CO2 + 2 H2O
1 CH4 + 2 O2  CO2 + 2 H2O
4 O present
on each side
When finally done, check for the
smallest coefficients possible

17. 17
Examples:
AgNO3 + H2S  Ag2S + HNO3
Al(OH)3 + H2SO4  Al2(SO4)3 + H2O
Fe3O4 + H2  Fe + H2O
C4H10 + O2  CO2 + H2O
2 AgNO3 + H2S  Ag2S + 2 HNO3
2 Al(OH)3 + 3 H2SO4  Al2(SO4)3 + 6 H2O
Fe3O4 + 4 H2  3 Fe + 4 H2O
2 C4H10 + 13 O2  8 CO2 + 10 H2O

18. 18
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

19. 19
SYNTHESIS or
COMBINATION
 In these reactions, 2 elements or compounds
combine to form another compound.
A + B  AB

20. 20
DECOMPOSITION
 In these reactions, a compound breaks up to form
2 elements or simpler compound.
AB  A +B

21. 21
SINGLE
REPLACEMENT
 In these reactions, a more reactive element
replaces a less reactive element in a compound.
A + BC  B + AC

22. 22
DOUBLE
REPLACEMENT
 In these reactions, two compounds combine to
form two new compounds.
AB + CD  AD + CB
 The cation from one compound replaces the cation
in another compound.
+ +

23. 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. 24
Examples:
Classify each of the reactions below:
1. Mg + CuCl2  MgCl2 + Cu
2. CaCO3  CaO + CO2
3. 2 HCl + Ca(OH)2  CaCl2 + 2 H2O
4. 4 Fe + 3 O2  2 Fe2O3
Single replacement
Mg is more reactive than Cu
Decomposition
Double replacement
Synthesis

25. Single Displacement
2
2 ZnCl
Cu(s)
(aq)
CuCl
Zn(s) 


The Zinc Replaces the Copper

26. 26
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.

27. 27
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)

28. 28
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

29. 29
Example 1:
Use activity series to complete each reaction below.
If no reaction occurs, write “No Reaction”.
Pb (s) + HCl (aq)  Metals
Fe
Ni
Sn
Pb
H
Cu
Ag
Pb is more
reactive
than H
PbCl2 (aq) + H2 (g)
2

30. 30
Example 2:
Use activity series to complete each reaction below.
If no reaction occurs, write “No Reaction”.
Ni (s) + CuCl2 (aq)  Metals
Fe
Ni
Sn
Pb
H
Cu
Ag
Ni is more
reactive
than Cu
NiCl2 (aq) + Cu (s)

31. 31
AQUEOUS
REACTIONS
 Many ionic solids dissolve in water to form ions.
2
H O
(s) (aq
2-
4 4 aq)
+
(
2 )
K 2
CrO CrO
K +


 These substances are called electrolytes.
2
H O
(s) (aq) (
-
3 2 3 q
2+
a )
Ba B
(NO ) 2 N
+ O
a


NaCl(s) Na+
(aq) + Cl
(aq)
H2O

32. 32
AQUEOUS
REACTIONS
 When ionic substances dissolve in water they
separate into ions.
K2CrO4 Ba(NO3)2

33. 33
 electrolytes are
substances whose
water solution is a
conductor of
electricity
 electrolytes are ions
dissolved in water
(Na+ + Cl-)
AQUEOUS
REACTIONS

34. 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

35. 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

36. 36
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

37. 37
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.
2 4 (aq) 3 2 (aq) 4 (s) 3 (aq)
K CrO + Ba(NO ) BaCrO + 2 KNO


Precipitate

38. Example of a Precipitation Reaction
Pb(NO3)2(aq) + 2 KI(aq)  2 KNO3(aq) + PbI2(s)

39. Let’s Look Closer at PbI2 Formation
Pb(NO3)2(aq) + 2 KI(aq)  2 KNO3(aq) + PbI2(s)

40. 40
SOLUBILITY
RULES
S
O
L
U
B
L
E
NO3
 No exceptions
Na+, K+
NH4
+
No exceptions
Cl, Br, I Except those containing Ag+
, Pb2+
SO4
2
Except those containing
Ba2+ , Pb2+, Ca2+
 Chemists use a set of solubility rules to predict
whether a product is soluble or insoluble.

41. 41
SOLUBILITY
RULES
I
N
S
O
L
S2, CO3
2
PO4
3
Except those containing
Na+ , K+, NH4
+
OH Except those containing
Na+ , K+, Ca2+, NH4
+

42. 42
Write balanced equations for each reactions shown
below. Indicate if no reaction occurs.
Example 1:
NaCl (aq) + AgNO3 (aq)  NaNO3 (?) + AgCl (?)
NaCl (aq) + AgNO3 (aq)  NaNO3 (aq) + AgCl (s)
soluble
precipitate

43. 43
Write balanced equations for each reactions shown
below. Indicate if no reaction occurs.
Example 2:
NH4Cl (aq) + KNO3 (aq)  NH4NO3 (?) + KCl (?)
NH4Cl (aq) + KNO3 (aq)  NH4NO3 (aq) + KCl (aq)
soluble
No Reaction

44. 44
Write balanced equations for each reactions shown
below. Indicate if no reaction occurs.
Example 3:
PbCl2 (aq) + NaI (aq)  PbI2 (?) + 2 NaCl (?)
precipitate
2
PbCl2 (aq) + 2 NaI (aq)  PbI2 (s) + 2 NaCl (aq)

45. © 2012 Pearson Education, Inc.
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.

46. © 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.

47. © 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.

48. © 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)

49. 49
NEUTRALIZATION
REACTIONS
 The most important reaction of acids and
bases is called neutralization.
 In these reactions an acid combines with a
base to form a salt and water.
2
HCl (aq) + NaOH (aq) NaCl (aq) + H O (l)
¾ ¾
®
Acid Base Salt
 Salts are ionic substances with the cation
donated from the base and the anion donated
from the acid.

50. 50
Write balanced equations for each of the neutral-
ization reactions shown below:
Examples:
HNO3 + Ba(OH)2  Ba(NO3)2 + H2O
2 2
H2SO4 + NaOH  Na2SO4 + H2O
2 2

51. 51
GAS FORMING
REACTIONS
 Some chemical reactions produce gas because
one of the products formed in the reaction is
unstable.
 Three such products are:
H2CO3 (aq)  CO2 (g) + H2O (l)
Carbonic acid:
Sulfurous acid: H2SO3 (aq)  SO2 (g) + H2O (l)
Ammonium: NH4OH (aq)  NH3 (g) + H2O (l)

52. 52
GAS FORMING
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)
2 HNO3 + K2SO3  2 KNO3 + H2SO3
2 HNO3 + Na2SO3  2 KNO3 + SO2 (g)+ H2O (l)

53. © 2012 Pearson Education, Inc.
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.

54. © 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.

55. © 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.

56. © 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.

57. 57
HEAT IN CHEMICAL
REACTIONS
 Reactions that release heat are classified as
exothermic.
 Reactions that absorb heat are classified as
endothermic.
H2 (g) + Cl2 (g)  2 HCl (g) + 185 kJ
 In exothermic reaction, heat is produced and can
be written as a product.
 In endothermic reaction, heat is required and
can be written as a reactant.
N2 (g) + O2 (g) + 181 kJ  2 NO (g)
Endothermic
Exothermic

58. © 2012 Pearson Education, Inc.
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.

59. © 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.

60. © 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.

61. © 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.

62. © 2012 Pearson Education, Inc.
Example 1:

63. Classifying Reactions
• Also we can classify reactions by what
happens:
• Redox reactions are the exchange of e-
• Redox are all reactions except?
Note

64. 64
OXIDATION-REDUCTION
REACTIONS
 Reactions known as oxidation and reduction
(redox) have many important applications in our
everyday lives.
 Rusting of a nail or the reaction within your car
batteries are two examples of redox reactions.
 In an oxidation-reduction reaction, electrons are
transferred from one substance to another.
 If one substance loses electrons, another substance
must gain electrons.

65. 65
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:
 Combination, decomposition, single
replacement and combustion reactions are
all examples of redox reactions.
Oxidation Is Loss of
electrons
Reduction Is Gain of
electrons
OIL
RIG

66. 66
OXIDATION-REDUCTION
REACTIONS
 In general, atoms of metals lose electrons to
form cations, and are therefore oxidized,
while atoms of non-metals gain electrons to
form anions, and are therefore reduced.
Ca + S CaS
Ca Ca2+ + 2
e-
S + 2 e- S2-
Oxidation
Reduction
 For example, in the formation of calcium
sulfide from calcium and sulfur
 Therefore, the formation of calcium sulfide
involves two half-reactions that occur
simultaneously, one an oxidation and the other
a reduction.

67. 67
COMBUSTION
 A reaction that involves oxygen as a reactant
and produces large amounts of heat is classified
as a combustion reaction.
 Combustion reactions are a subclass of
Oxidation-Reduction reactions
occurs in
the cylinders
of the engine
2 C8H18(g) + 25 O2(g)  16 CO2(g) + 18 H2O(g)

68. 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)

69. 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.

70. 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

71. 71
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
loss of oxygen or gain of hydrogen.
 For example, poisonous methyl alcohol is
metabolized by the body by the following reaction:
CH3OH H2CO + 2H•
methyl alcohol formaldehyde
Oxidation
(loss of
hydrogen)

72. 72
REDOX IN
BIOLOGICAL SYSTEMS
 The formaldehyde is further oxidized to formic
acid and finally carbon dioxide and water by the
following reactions:
2 H2CO + O2 2H2CO2
formic acid
formaldehyde
Oxidation
(gain of
oxygen)
formic acid
2 H2CO2 + O2 CO2 + H2O

73. 73
REDOX IN
BIOLOGICAL SYSTEMS
 In many biochemical oxidation-reduction reactions,
the transfer of hydrogen atoms produces energy in
the cells.
 The oxidation of a typical biochemical molecule can
involve the transfer of two hydrogen atoms to a
proton acceptor such as coenzyme FAD to produce
its reduced form FADH2.

74. 74
REDOX IN
BIOLOGICAL SYSTEMS
 In summary, the particular definition of oxidation-
reduction depends on the process that occurs in the
reaction.

75. 75
THE END