Download Chapter 8

Chemical Equations
Chapter 8
Hein and Arena
Version 2.0
12th Edition
Eugene Passer
Chemistry Department
Bronx Community
1 College
© John Wiley and Sons, Inc
Chapter Outline
8.1 The Chemical Equation
8.4 Types of Chemical Equations
8.2 Writing and Balancing
Chemical Equations
8.5 Heat in Chemical Reactions
8.3 Information in a
Chemical Equation
2
• Chemists use chemical equations to
describe reactions they observe in the
laboratory or in nature.
• Chemical equations provide us with
the means to
1. summarize the reaction
2. display the substances that are reacting
3. show the products
4. indicate the amounts of all component
substances in a reaction.
3
8.1
The Chemical Equation
4
• Chemical reactions always involve
change.
• Atoms, molecules or ions rearrange to
form different substances.
• The substances entering the reaction
are called reactants.
• The substances formed in the reaction
are called products.
• During reactions, chemical bonds are
broken and new bonds are formed.
5
• A chemical equation is a shorthand
expression for a chemical change or
reaction.
• A chemical equation uses the chemical
symbols and formulas of the reactants
and products and other symbolic terms
to represent a chemical reaction.
6
Chemical Equation
Al + Fe2O3  Fe + Al2O3
reactants
products
iron oxygen
bonds break
aluminum oxygen
bonds form
7
Coefficients (whole numbers) are placed
in front of substances to balance the
equation and to indicate the number of
units (atoms, molecules, moles, or ions)
of each substance that are reacting.
8
2 Al + Fe2O3  2 Fe + Al2O3
coefficient
coefficient
9
Conditions required to carry out the
reaction may be placed above or below
the arrow.
10

2 Al + Fe2O3  2 Fe + Al2O3
coefficient
coefficient
 heat
11
The physical state of a substance is
indicated by symbols such as (l) for
liquid.
12
In a chemical
reaction atoms are
neither created nor
destroyed.
2Al(s)
(s) + Fe2O3(s)  2Fe(l)
(s) + Al2O3 (s)
All atoms present in
the reactant must
also be present in
the products.
13
Symbols Used
in Chemical Reactions
14
symbol
+
meaning
plus
location
placed
between
substances
15
symbol

meaning
yields
location
between
reactants and
products
16
symbol
(s)
meaning
solid
location
after formula
17
symbol
(l)
meaning
liquid
location
after formula
18
symbol
(g)
meaning
gas
location
after formula
19
symbol
(aq)
meaning
aqueous
location
after formula
20
symbol

meaning
heat
location
written above or
below 
21
symbol
h
meaning
light energy
location
written above 
22
symbol

meaning
gas formation
location
after formula
23
8.2
Writing and
Balancing Equations
24
• To balance an equation adjust the
number of atoms of each element so
that they are the same on each side of
the equation.
• Never change a correct formula to
balance an equation.
25
Steps for
Balancing Equations
26
Step 1 Identify the reaction. Write a
description or word equation for the
reaction.
Mercury (II) oxide decomposes to form mercury
and oxygen.
mercury(II) oxide → mercury + oxygen
27
Step 2 Write the unbalanced (skeleton)
equation.
– The formulas of the reactants and
products must be correct.
– The reactants are written to the left of
the arrow and the products to the right
of the arrow.
HgO  Hg + O2
The formulas of the reactants and
products can never be changed.
28
Step 3a Balance the equation.
– Count and compare the number of atoms of each
element on both sides of the equation.
– Determine the elements that require balancing.
29
Step 3a Balance the equation.
HgO → Hg + O2
Element
Hg
Reactant Side
1
Product Side
1
– There is one mercury atom on the reactant side
and one mercury atom on the product side.
– Mercury is balanced.
30
Step 3a Balance the equation.
HgO  Hg + O2
Element
O
Reactant Side
1
Product Side
2
– There are two oxygen atoms on the product
side and there is one oxygen atom on the
reactant side.
– Oxygen needs to be balanced.
31
Step 3b Balance the equation.
– Balance each element one at a time, by
placing whole numbers (coefficients) in front
of the formulas containing the unbalanced
element.
– A coefficient placed before a formula
multiplies every atom in the formula by that
coefficient.
32
Step 3b Balance the equation.
2HgO  Hg + O2
Element
O
Reactant Side
12
Product Side
2
– Place a 2 in front of HgO to balance O.
– There are two oxygen atoms on the reactant side
and there are two oxygen atoms on the product
side.
– Oxygen (O) is balanced.
33
Step 3c Balance the equation.
– Check all other elements after each individual
element is balanced to see whether, in
balancing one element,
another element
became unbalanced.
34
Step 3c Balance the equation.
2HgO  Hg + O2
Element
Hg
Reactant Side
2
Product Side
1
– Count and compare the number of mercury
(Hg) atoms on both sides of the equation.
– There are two mercury atoms on the
reactant side and there is one mercury
atom on the product side.
– Mercury (Hg) is not balanced.
35
Step 3c Balance the equation.
2HgO  2Hg + O2
Element
Hg
Reactant Side
2
Product Side
12
– Place a 2 in front of Hg to balance
mercury.
– There are two mercury atoms on the
reactant side and there are two mercury
atoms on the product side.
– Mercury (Hg) is balanced.
36
 THE EQUATION IS BALANCED
2HgO  2Hg + O2
Element
Hg
O
Reactant Side
2
2
Product Side
2
2
37
Balance the Equation
sulfuric acid + sodium hydroxide → sodium sulfate + water
38
Balance the Equation
H2SO4(aq) + 2NaOH(aq) → Na2SO4(aq) + H2O(l)
Reactant Side
1
Product Side
1
Na
12
2
O
12
1
H
34
2
SO
24
Place ais2one
There
in front
Na on
ofthe
NaOH
reactant
to balance
side and
Na.there are
two Na on the product side.
39
 THE EQUATION IS BALANCED 
2 +
H2SO4(aq)2 + NaOH(aq) → Na2SO4(aq)
H2O(l)
Reactant Side
Product Side
21
1
SO 4
Na
2
2
O
2
12
H
4
42
There
4H
on the
side and two H on
Place aare
2 in
front
of Hreactant
2O to balance H.
the product side.
40
Balance the Equation
butane + oxygen → carbon dioxide + water
41
Balance the Equation
C4H10 (g) + O2 (g) → 4 CO2(g) + H2O(l)
C
Reactant Side
4
Product Side
14
H
10
2
O
2
39
Place aare
There
4 in
four
front
C on
of the
CO2reactant
to balance
sideC.
and there is
one C on the product side.
42
C4H10 (g) + O2 (g) → 4 CO2(g) + 5H2O(l)
C
Reactant Side
4
Product Side
4
H
10
2
10
O
2
9
13
Place aare
There
5 in
10front
H onofthe
H2reactant
O to balance
side and
H. there are
two H on the product side.
43
10
2 C4H10 (g) + O2 (g) → 48 CO2(g) +10
55 H2O(l)
C
Reactant Side
48
Product Side
84
H
10
20
10
20
O
2
13
26
To balance
There
is no O,
whole
double
number
the coefficients
coefficient of
that
each
can be
placed in front
substance
otherofthan
O2 oxygen.
to balance O.
44
 THE EQUATION IS BALANCED 
2C4H10 (g) +13O2 (g) → 8CO2(g) +10H2O(l)
C
Reactant Side
8
Product Side
8
H
20
20
O
2
26
26
There aare
Place
13now
in front
26 Oofon
O2the
to product
balance side.
O.
45
8.3
Information in a
Chemical Equation
46
The meaning of a formula
is context dependent.
The formula H2O can mean:
1.
2.
3.
4.
5.
2H and 1 O atom
1 molecule of water
1 mol of water
6.022 x 1023 molecules of water
18.02 g of water
47
In an equation, formulas can represent units
of individual chemical entities or moles.
H2
+
1 molecule H2
1 mol H2
Cl2
1 molecule Cl2
1 mol Cl2
→
2HCl
2 molecules HCl
2 mol HCl
48
Formulas
Number of
molecules
Number
of atoms
Number
of
moles
Molar
masses
49
8.4
Types of Chemical
Equations
50
Combination
Decomposition
Single-Displacement
Double-Displacement
51
Combination Reactions
52
Two reactants combine to form one product.
A + B  AB
53
Examples
54
Metal + Oxygen → Metal Oxide
2Ca(s) + O2(g)  2CaO(s)
4Al(s) + 3O2(g)  2Al2O3(s)
55
Nonmetal + Oxygen → Nonmetal Oxide
S(s) + O2(g)  SO2(g)
N2(g) + O2(g)  2NO(g)
56
Metal + Nonmetal → Salt
2K(s) + F2(g)  2KF(s)
2Al(s) + 3Cl2(g)  2AlCl3(s)
57
Metal Oxide + Water → Metal Hydroxide
Na2O(s) + H2O(l)  2NaOH(aq)
CaO(s) + 2H2O(l)  2Ca(OH)2(aq)
58
Nonmetal Oxide + H2O(l) → Oxy-acid
SO3(g) + H2O(l)  H2SO4(aq)
N2O5(g) + H2O(l)  2HNO3(aq)
59
Decomposition Reactions
60
A single substance breaks down to
give two or more different substances.
AB  A + B
61
Examples
62
Metal Oxide → Metal + Oxygen
2Ag2O(s)  4Ag(s) + O2(g)
Metal Oxide → Metal Oxide + Oxygen
2PbO2(s)  2PbO(s) + O2(g)
63
Carbonate → CO2(g)
CaCO3(s)  CaO(s) + CO2(g)
Hydrogen Carbonate → CO2(g)
2NaHCO3(s)  Na2CO3(s) + H2O(g) + CO2(g)
64
Miscellaneous Reactions
2H2O2(l)  2H2O(l) + O2(g)
2KClO3(s)  2KCl(s) + 3O2(g)
2NaNO3(s)  2NaNO2(s) + O2(g)
65
Single Displacement
Reactions
66
One element reacts with a compound to
replace one of the elements of that compound.
A + BC  AC + B
67
Metal + Acid → Hydrogen + Salt
Mg(s) + 2HCl(aq)  H2(g) + MgCl2(aq)
salt
2Al(s) + 3H2SO4(aq)  3H2(g) + Al2(SO4)3(aq)
salt
68
Metal + Water → Hydrogen + Metal Hydroxide
2Na(s) + 2H2O(l)  H2(g) + 2NaOH(aq)
metal
hydroxide
Ca(s) + 2H2O(l)  H2(g) + Ca(OH)2(aq)
metal
hydroxide
69
Metal + Water → Hydrogen + Metal Oxide
3Fe(s) + 4H2O(g)  4H2(g) + Fe3O4(s)
metal
oxide
70
The Activity Series
71
An atom of an element in the activity series will displace an
atom of an element below it from one of its compounds .
increasing
activity
Metals
K
Ca
Na
Mg
Al
Zn
Sodium (Na) will displace an
Fe
atom below it from one of its
Ni
Sn
compounds.
Pb
H
Cu
Ag
Hg
72
Examples
Metal Activity Series
73
Metal Higher in Activity Series Displacing Metal Below It
Mg(s) + PbS(s)  MgS(s) + Pb(s)
Metals
Mg
Al
Zn
Fe
Ni
Sn
Pb
Magnesium is above lead
in the activity series.
74
Metal Lower in Activity Cannot Displace Metal Above It
Ag(s) + CuCl2(s)  no reaction
Metals
Pb
H
Cu
Ag
Hg
Silver is below copper in
the activity series.
75
Example
Halogen Activity Series
76
Halogen Higher in Activity Series Displaces Halogen Below It
Cl2(g) + CaBr2(s)  CaCl2(aq) + Br2(aq)
Halogens
F2
Cl2
Br2
I2
Chlorine is above bromine
in the activity series.
77
Double Displacement
Reactions
78
Two
compounds
exchange
each
The reaction
can be
thoughtpartners
of as anwith
exchange
A
B displaces C
D and combines with D
C
other
to produce
two different
compounds.
of positive
and negative
groups.
AB + CD  AD + CB
79
The Following Accompany Double
Displacement Reactions
•
•
•
•
formation of a precipitate
release of gas bubbles
release of heat
formation of water
80
Examples
81
Acid Base Neutralization
acid + base → salt + water
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
H2SO4(aq) + 2NaOH(aq)  Na2SO4(aq) + 2H2O(l)
82
Formation of an Insoluble Precipitate
AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)
Pb(NO3)2(aq) + 2KI(aq)  PbI2(s) + 2KNO3(aq)
83
Metal Oxide + Acid
metal oxide + acid → salt + water
CuO(s) + 2HNO3(aq)  Cu(NO3)2(aq) + H2O(l)
CaO(s) + 2HCl(aq)  CaCl2(s) + H2O(l)
84
Formation of a Gas
H2SO4(aq) + 2NaCN(aq)  Na2SO4(aq) + 2HCN(g)
NH4Cl(aq) + NaOH(aq)  NaCl(aq) + NH4OH(aq)
indirect gas
formation
NH4OH(aq)  NH3(g) + H2O(l)
85
8.5
Heat in
Chemical Reactions
86
Energy changes always accompany
chemical reactions.
Whenreason
One
this why
occurs,
reactions
energy
occur
is
is that thetoproduct
released
the surroundings.
attains a lower
energy state than the reactants.
87
The amounts
of substances
areliberate
expressed
in moles.
Exothermic
reactions
heat.
H2(g) + Cl2(g) → 2HCl(g) + 185 kJ (exothermic)
1 mol 1 mol
2 mol
Endothermic reactions absorb heat.
N2(g) + O2(g) + 185 kJ → 2NO(g) (exothermic)
1 mol 1 mol
2 mol
88
For life on Earth the sun is
the major provider of energy.
The energy for plant photosynthesis is
derived from the sun.
6CO2 + 6H2O + 2519 kJ → C6H12O6 + 6O2
glucose
glucose
89
Energy of Activation
90
• A certain amount of energy is always
required for a reaction to occur.
• The energy required to start a reaction
is called the energy of activation.
91
CH4 + 2O2 → CO2 + 2H2O + 890 kJ
• This reaction will not occur unless
activation energy is supplied.
• The activation energy can take the
form of a spark or a flame.
92
93
8.2
8.3
94