Chapter 7
Chemical Reactions
Homework
► Assigned
Problems (odd numbers only)
► “Questions
and Problems” 7.1 to 7.31 (begins
on page 200)
► “Additional
Questions and Problems” 7.41 to
7.49 (page 221)
► “Challenge
Questions” 7.51-7.57 (page 222)
Chemical Reactions
► Physical
changes:
 Involves no changes in chemical identity of a
substance
 No changes in physical properties (color,
physical state, freezing point, boiling point)
► Chemical
changes:
 A chemical reaction in which one or more
substances changes to a different substance
 Properties that matter exhibits as it undergoes
changes in chemical composition
Chemical Reactions
► Chemical
properties determine whether or
not a substance can be changed to another
substance
► Reactions involve chemical changes in
matter resulting in new substances
► Reactions involve rearrangement and
exchange of atoms to produce new
molecules
 Elements are not changed during a reaction
Reactants  Products
►A
Changes During
Chemical Reactions
chemical change occurs when new
substances are made
 Conversion of material(s) into one or more
new substances
 These substances will have different properties
from the original material
►New
properties are visible (visual
clues)
 Color change, precipitate formation, gas
bubbles, flames, heat release
Changes During Chemical Reactions
Fe
Fe2O3
Li
LiOH, H2
HCO3Na
CO2
NaOH, H2
Changes During Chemical Reactions
►In a chemical reaction:
 At least one new substance is
produced
 Atoms are never created or
destroyed
 Every atom present as a reactant has
to be present as a product
 The atoms in reactants rearrange to
form new products
Chemical Equations
►A
chemical equation is a written statement
that uses symbols and formulas (no words)
to describe the changes during a chemical
reaction
► It shows substances at the beginning of a
reaction (reactants)
► It shows substances formed in the reaction
(products)
Writing a Chemical Equation
► Chemical
reactions can be written as:
 Word equations
 Formula equations
reactants
products
Balancing Chemical Equations
►A
balanced chemical reaction has the same
number of atoms of each element on both
sides of the arrow
► Atoms are neither created nor destroyed
► Every atom must be accounted for
► Equations are balanced by placing a
coefficient in front one or more of the
substances in the equation
Symbols Used in Equations
►Symbols
used after chemical
formula to indicate physical state
 (g) = gas
 (l) = liquid
 (s) = solid
 (aq) = aqueous, dissolved in water
Writing Chemical Equations
When magnesium metal burns in air it
produces a white, powdery compound
magnesium oxide
►

Burning in air means reacting with O2
Write the word equation
►



The reactants are to the left of the arrow
The products are to the right of the arrow
Two or more reactants or products are separated by
a plus sign
magnesium + oxygen
magnesium oxide
Writing Chemical Equations
► Indicate
the physical state of each substance
 Use the correct chemical symbol to indicate
liquids and solids
 Metals are solids, except for Hg which is liquid
 Use molecular form for gases (H2, O2, N2, all
halogens)
 Identify polyatomic ions
magnesium(s) + oxygen(g)
magnesium oxide(s)
Writing Chemical Equations
► Convert
the word equation into a formula
equation
 Use the correct chemical symbol to indicate
liquids and solids
 There must be the same number of each kind
of atom on the reactant and product side of the
equation
___Mg (s) +___O2 (g)
___MgO(s)
 Determine if the equation is balanced
 If not equal, must BALANCE
Balancing Chemical Equations
►Balance
equations by the use of a
coefficient placed to the left of a
substance
►NEVER change the subscripts of a
compound to balance an element
 It changes the identity of the compound
►Can
change coefficients but never
subscript numbers
Balancing a Chemical Equation
Example 1
2 Mg(s) 
O2 (g)  2 MgO(s)
Coefficient
2
1 Mg
2O
2 1 Mg
2 1O
Balancing a Chemical Equation:
Example 2
►When
solid ammonium nitrite is
heated it produces nitrogen gas
and water vapor
 Write the formula equation
NH4 NO2 (s) 
N2 (g) 
H2 O(g)
Balancing a Chemical Equation
Example 2
NH4 NO2 (s) 
2xN
2xO
4xH
N2 (g)  2 H2 O(g)
2xN
21 x O
42 x H
Balancing a Chemical Equation:
Example 3
►Nitrogen
monoxide gas decomposes to
produce dinitrogen monoxide gas and
nitrogen dioxide gas
 Write the formula equation
NO(g) 
N2 O(g) 
NO2 (g)
Balancing a Chemical Equation
Example 3
3 NO(g)  N2O (g)  NO2 (g)
31 x N
31 x O
3xN
3xO
Balancing a Chemical Equation
Example 4
►Liquid nitric acid decomposes to
reddish-brown nitrogen dioxide gas,
liquid water and oxygen gas.
 Write the formula equation
HNO3 (l) 
NO 2 (g) 
H 2 O(l) + O2 (g)
Balancing a Chemical Equation
Example 4
4
4
2HNO3 (l)  2NO 2 (g)  2H 2 O(l) + O2 (g)
421 x N
12 6 3 x O
42 1 x H
42 1 x N
12 7 5 x O
42 x H
Types of Reactions
► Reactions
are separated into groups of
similar reactions
► Based on the form of the equation for the
reaction





Synthesis (combination)
Decomposition
Single replacement
Double replacement
Combustion
Types of Reactions
►Synthesis
Reactions
 Reactions in which two or more
substances combine to form a third
substance
 (one product forms)
►General
form of equation:
A+B
AB
Synthesis Reactions
► The
combinations can include
 Two elements
 An element and a compound
 Two compounds
► Examples
2Na (s)  Cl2 (g)  2 NaCl (s)
2NO (g)  O2 (g)  2 NO2 (g)
NH3 (g)  HCl (g)  NH4Cl(s)
Types of Reactions
► Decomposition
Reactions
 Reactions in which one reactant breaks down
into simpler (smaller) substances
 Generally initiated by addition of energy
(electric current or heating substances to high
temperature)
 Opposite of a Synthesis Reaction
► General
Form of Equation
AB
A+B
Decomposition Reactions
► Can
be broken down to:
 Smaller compounds
 Elements
 Both
► Examples

CaCO 3 (s) 
 CaO (s)  CO 2 (g)
2 AgBr (s) light

 2 Ag (s)  Br ()
2

2 N2O5 (g) 

4 NO 2 (g)  O2 (g)
Types of Reactions
► Single
replacement reactions
 One element replaces another element
 Forms a new compound which frees the
replaced element
 Most reactions occur in an aqueous solution
► General
Form of Equation
A + BC
AC + B
Single Replacement Reactions
► Three
types
 Metal replaces a metal
 Metal replaces hydrogen
 Nonmetal replaces nonmetal
► Examples
Cu (s)  2 AgNO3 (aq)  Cu(NO 3 )2 (aq)  2 Ag (s)
metal
replaces
metal
Zn (s)  H2 SO 4 (aq)  ZnSO4 (s)  H2 (g)
metal
replaces
hydrogen
2 NaI (aq)  Cl 2 (aq)  2 NaCl (aq)  I2 (aq)
nonmetal
replaces
nonmetal
Type of Reactions
► Two
compounds exchange ions or atoms to
form new compounds
► Also called exchange reactions
 Shows the exchange of “associates” when
comparing the reactants and products
► General
Form of Equation
AB + CD
AD + BC
Double Replacement Reactions
► Most
of these reactions occur in aqueous
solution
► Most involve acids, bases, and ionic
compounds
► Products formed
 Precipitate (a solid that is insoluble)
 A gas
 Water
Double Replacement Reactions
► Examples
AgNO 3 (aq)  NaCl (aq)  NaNO 3 (aq)  AgCl (s)
precipitate
Na2S (aq)  2HCl (aq)  2NaCl (aq)  H2S (g)
gas
HCl (aq)  NaOH (aq)  NaCl (aq)  H2O ()
water
Summary of Reaction Types
Combustion Reactions
► Occurs
when a hydrocarbon combines with
oxygen which produces carbon dioxide,
water and heat (flame)
► The reaction of oxygen with any substance
► If a combustion reaction is possible then the
substance will burn
Combustion Reactions
► Examples
 The combustion of propane gas
►Produces carbon dioxide and water
►Produces heat (flame)
C3H8 (g)  5 O 2 (g)  3CO2 (g)  4 H2 O (g)
hydrocarbon
 The combustion of sulfur
►Also a combination reaction
►Also produces heat (flame)
S (s)  O2 (g)  SO2 (g)
Energy in Chemical Reactions
► In
a chemical reaction
 A change in energy occurs as bonds are broken
(reactants) and new ones form (products)
 Nearly all chemical reactions absorb or produce
heat
 Measured by the heat of reaction or enthalpy
► Enthalpy
change is the amount of heat
produced or consumed in a process (∆H )
Heat of Reaction
► Endothermic
occur
reactions absorb heat as they
 If (∆H ) is positive, then heat is added to the
reaction
► Exothermic
occur
reactions produce heat as they
 If (∆H ) is negative, then heat is evolved by the
reaction
Heat of Reaction
► Photosynthesis
reaction
 Carbon dioxide reacts with water to produce glucose
and oxygen
6 CO2 (g)  6 H2 O ()  C6H12 O 6 (s)  6 O 2 (g)
∆H = +2801 kJ
► Cell
metabolism
 Glucose reacts with oxygen to produce carbon dioxide
and water
C6H12 O 6 (s)  6 O 2 (g)  6 CO2 (g)  6 H2 O ()
∆H = -2801 kJ
Calculation of Heat in Reactions
► The
combustion of sulfur dioxide
► It reacts with oxygen to produce sulfur trioxide
2 SO2 (g)  O2 (g)  2 SO3 (g)
∆H = -99.1 kJ
► Calculate
the heat produced when 75.2 g of
sulfur trioxide is produced
Given 75.2 g SO3
Heat in kJ produced
when SO3 is formed
Calculation of Heat in Reactions
Relation between g of SO3 and heat released
Grams
of SO3
Molar
mass
Moles
of SO3
Heat
of rxn
kj
Write the necessary conversion factors
1 mol SO3
80.07 g
and
80.07 g
1 mol SO3
2 mol SO3
99.1 kJ
and
99.1 kJ
2 mol SO3
Set up the problem
1mol SO3
99.1kJ
75.2 g SO3 

 46.5 kJ
80.07 g SO3 2 mol SO3
►End