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Resources
Bellringers
Chapter Presentation
Transparencies
Standardized Test Prep
Math Skills
Visual Concepts
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Chapter 6
The Nature of Chemical Reactions
Table of Contents
Section 1 The Nature of Chemical Reactions
Section 2 Reaction Types
Section 3 Balancing Chemical Equations
Section 4 Rates of Change
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Chapter 6
Section 1 The Nature of
Chemical Reactions
Objectives
• Recognize some signs that a chemical reaction may
be taking place.
• Explain chemical changes in terms of the structure
and motion of atoms and molecules.
• Describe the differences between exothermic and
endothermic reactions.
• Identify situations involving chemical energy.
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Chapter 6
Section 1 The Nature of
Chemical Reactions
Bellringer
Methane, CH4, is an organic compound that is the
principal component of natural gas. Many people burn
methane when cooking or heating homes. The chemical
reaction of methane burning is shown in several
ways below.
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Chapter 6
Section 1 The Nature of
Chemical Reactions
Bellringer
1. What else besides carbon dioxide and water is produced in this
reaction that makes methane useful
for cooking and heating?
2. Complete the table below with the number of atoms of each
element before and after the reaction.
3. How does the number of atoms of each element on the left
side of the equation compare to the number on the right? What
law does this demonstrate?
4. Use your answer to item 1 and the law of conservation of
energy to guess whether there is more energy stored in the
bonds among the atoms before the reaction or among the
bonds of the atoms after the reaction.
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Chapter 6
Section 1 The Nature of
Chemical Reactions
Chemical Reactions Change Substances
• Chemical reactions occur when substances undergo
chemical changes to form new substances.
•
•
•
•
•
Signs of chemical reactions are:
Production of gas
Formation of a solid
change of color
Energy released or absorbed
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Chapter 6
Section 1 The Nature of
Chemical Reactions
Chemical Reactions Change Substances
• Chemical reactions rearrange atoms.
• A reactant is a substance or molecule that
participates in a chemical reaction.
• A product is a substance that forms in a
chemical reaction.
• CH4 + 2O2  2H2O + CO2 + Energy
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Chapter 6
Section 1 The Nature of
Chemical Reactions
Chemical Reaction
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Chapter 6
Section 1 The Nature of
Chemical Reactions
Signs of a Chemical Reaction
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Chapter 6
Section 1 The Nature of
Chemical Reactions
Energy and Reactions
• Energy must be added to break bonds.
• Many forms of energy can be used to break bonds:
•
•
•
•
heat
electricity
sound
light
• Forming bonds releases energy.
• Example: When gasoline burns, energy in the form
of heat and light is released as the products of the
isooctane-oxygen reaction and other gasoline
reactions form.
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Chapter 6
Section 1 The Nature of
Chemical Reactions
Reaction Model
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Chapter 6
Section 1 The Nature of
Chemical Reactions
Energy and Reactions, continued
• Energy is conserved in chemical reactions.
• Chemical energy is the energy released when a
chemical compound reacts to produce new compounds.
• The total energy that exists before the reaction is equal to
the total energy of the products and their surroundings.
• An exothermic reaction is a chemical reaction in
which heat is released to the surroundings.
• An endothermic reaction is a chemical reaction
that absorbs heat.
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Chapter 6
Section 1 The Nature of
Chemical Reactions
Energy and Reactions, continued
• The graphs below represent the changes in chemical
energy for an exothermic reaction and an
endothermic reaction.
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Chapter 6
Section 1 The Nature of
Chemical Reactions
Objectives Review
• Recognize some signs that a chemical reaction may
be taking place.
• Explain chemical changes in terms of the structure
and motion of atoms and molecules.
• Describe the differences between exothermic and
endothermic reactions.
• Identify situations involving chemical energy.
• Section 1 Review Page 189 # 1-6
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Chapter 6
Section 2 Reaction Types
Objectives
• Distinguish among five general types of chemical
reactions.
• Predict the products of some reactions based on the
reaction type.
• Describe reactions that transfer or share electrons
between molecules, atoms, or ions.
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Chapter 6
Section 2 Reaction Types
Bellringer
There are thousands of ways that more than one hundred
elements can combine with each other to form different
substances. Just as the elements can be sorted into
families, the many reactions the elements undergo can be
classified as a few basic types. The types of reactions are
classified based on whether they involve combining atoms
or smaller molecules to make larger molecules (synthesis),
breaking down larger molecules into atoms or smaller
molecules (decomposition), or having atoms of one
element replace the atoms of another element within a
compound (single- or double-displacement).
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Chapter 6
Section 2 Reaction Types
Bellringer, continued
1. In which reaction model do three elements combine to make a compound?
2. In which reaction model is a complex substance broken down into simpler
parts?
3. Identify the reaction model in which one element reacts with a compound,
leaving behind another element and a new compound containing the first
element.
4. In which reaction model do two compounds react to form two
different compounds?
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Chapter 6
Section 2 Reaction Types
Classifying Reactions
• A synthesis reaction is a reaction in which two or
more substances combine to form a new compound.
• Synthesis reactions have the following general
form: A + B → AB
• Example: In the following synthesis reaction, the
metal sodium reacts with chlorine gas to form
sodium chloride, or table salt.
• 2Na + Cl2 → 2NaCl
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Chapter 6
Section 2 Reaction Types
Synthesis Reaction
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Chapter 6
Section 2 Reaction Types
Classifying Reactions, continued
• A decomposition reaction is a reaction in a single
compound breaks down to form two or more simpler
substances.
• Decomposition reactions have the following
general form: AB → A + B
• Example: The following shows the decomposition
of water.
• 2H2O → 2H2 + O2
• Electrolysis is the process in which an electric
current is used to produce a chemical reaction,
such as the decomposition of water.
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Chapter 6
Section 2 Reaction Types
Decomposition Reaction
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Chapter 6
Section 2 Reaction Types
Electrolysis
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Chapter 6
Section 2 Reaction Types
Classifying Reactions, continued
• A combustion reaction is the oxidation reaction of
an organic compound, in which heat is released.
• Combustion reactions use oxygen as a reactant
and release energy.
• CH4 + 2O2  2H2O + CO2 + Energy
• H2O and CO2 are common products of combustion
• In combustion the products depend on the amount of
oxygen available for the reaction.
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Chapter 6
Section 2 Reaction Types
Combustion Reaction
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Chapter 6
Section 2 Reaction Types
Classifying Reactions, continued
• A single-displacement reaction is a reaction in
which one element or radical takes the place of
another element or radical in the compound.
• Single-displacement reactions have the following
general form: AX + B → BX + A
• Example: The single-displacement reaction
between copper(II) chloride and aluminum is
shown as follows.
3CuCl2 + 2Al → 2AlCl3 + 3Cu
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Chapter 6
Section 2 Reaction Types
Single Displacement
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Chapter 6
Section 2 Reaction Types
Classifying Reactions, continued
• A double-displacement reaction is a reaction in
which a gas, a solid precipitate, or a molecular
compound forms from the apparent exchange of
atoms or ions between two compounds.
• Double-displacement reactions have the following
general form: AX + BY → AY + BX
• Example: The double-displacement reaction that
forms lead chromate is as follows.
Pb(NO3)2 + K2CrO4 → PbCrO4 + 2KNO3
Oxygen and Potassium went on a date…
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Chapter 6
Section 2 Reaction Types
Double Displacement Reaction
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Chapter 6
Section 2 Reaction Types
Electrons and Chemical Reactions
• An oxidation-reduction reaction is any chemical
change in which one species gains electrons and
another species loses electrons.
• Oxidation-reduction reactions are often called redox reactions
for short.
• Substances that accept electrons in a redox reaction are said
to be reduced.
• Substances that give up electrons in a redox reaction are said
to be oxidized.
• A radical is an organic group that has one or more
electrons available for bonding.
• Polymerization reactions can occur when radicals
are formed.
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Chapter 6
Section 2 Reaction Types
Redox Reactions
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Chapter 6
Section 2 Reaction Types
Objectives
• Distinguish among five general types of chemical
reactions.
• Predict the products of some reactions based on the
reaction type.
• Describe reactions that transfer or share electrons
between molecules, atoms, or ions.
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Chapter 6
Section 3 Balancing Chemical
Equations
Reaction #
Reactant #1
Reactant #2
1) name
Sodium
Phosphate
Iron (III)
Nitrate
1) formula
Na3PO4
Fe(NO3)3
Product #1
Product #2
Rx Evidence
2) name
2) formula
3) name
3) formula
4) name
4) formula
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Chapter 6
Section 3 Balancing Chemical
Equations
Objectives
• Demonstrate how to balance chemical equations.
• Interpret chemical equations to determine the
relative number of moles of reactants needed and
moles of products formed.
• Explain how the law of definite proportions allows for
predictions about reaction amounts.
• Identify mole ratios in a balanced chemical equation.
• Calculate the relative masses of reactants and
products from a chemical equation.
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Chapter 6
Section 3 Balancing Chemical
Equations
Bellringer
You have already used scientific shorthand by writing
symbols for elements and formulas for compounds. You
can use these formulas to write chemical equations that
summarize what happens during a chemical reaction
and how much of each substance is involved. Examine
the reaction model for the water synthesis reaction
shown on the next slide, and answer the items
that follow.
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Chapter 6
Section 3 Balancing Chemical
Equations
1. What is the difference between reaction models A and B?
2. Why is reaction model A not fully complete as written? (Hint: Consider
how many atoms of each element exist before and after the reaction.)
3. A friend tells you that an easier way to make sure the same number of
atoms are on both sides of the equation is to change the subscript on the
product so that it is H2O2 instead of H2O. What’s wrong with this
reasoning? (Hint: If you did this, would it still be a synthesis reaction
for water?)
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Chapter 6
Section 3 Balancing Chemical
Equations
Describing Reactions
• One way to record the products and reactants of a reaction is to
write a word equation.
• Example: methane + oxygen → carbon dioxide + water
• A chemical equation is a representation of a chemical reaction
that uses symbols to show the relationship between the reactants
and the products.
• In a chemical equation, such as the one above, the reactants,
which are on the left-hand side of the arrow, form the products,
which are on the right-hand side.
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Chapter 6
Section 3 Balancing Chemical
Equations
Describing Reactions
• When the number of atoms of reactants matches the
number of atoms of products, then the chemical
equation is said to be balanced.
• Balancing equations follows the law of
conservation of mass.
• You cannot balance chemical equations by changing
chemical formulas themselves, because that would
change the substances involved.
• To balance chemical equations, numbers called
coefficients must be placed in front of the
chemical formulas.
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Chapter 6
Section 3 Balancing Chemical
Equations
Law of Conservation of Mass
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Chapter 6
Section 3 Balancing Chemical
Equations
Describing Reactions, continued
• When the numbers of atoms for each element are the
same on each side, the equation is balanced, as
shown below.
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Chapter 6
Section 3 Balancing Chemical
Equations
Reading a Chemical Equation
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Chapter 6
Section 3 Balancing Chemical
Equations
Chemical Equation
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Chapter 6
Section 3 Balancing Chemical
Equations
Balancing a Chemical Equation by Inspection
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Balance by changing the Coefficients
___N2 + ___H2  ___ NH3
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Balance by changing the Coefficients
___O2  ___ O3
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Balance by changing the Coefficients
___K + ___H2O  ___H2 + ___KOH
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Balance by changing the Coefficients
___CH4 + ___O2  ___ CO2 + ___H2O
___C3H8 + ___O2  ___ CO2 + ___H2O
___HF + ___SiO2  ___ SiF4 + ___H2O
___NH4NO2  ___N2 + ___H2O
___NO  ___ N2O + ___NO2
___HNO3  ___NO2 + ___ H2O + ___O2
___NH3 + ___O2  ___ NO + ___H2O
___C2H5OH + ___O2  ___ CO2 + ___H2O
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Chapter 6
Section 3 Balancing Chemical
Equations
Math Skills
Balancing Chemical Equations Write the equation
that describes the burning of magnesium in air to
form magnesium oxide.
1. Identify the reactants and products.
Magnesium and oxygen gas are the reactants that
form the product, magnesium oxide.
2. Write a word equation for the reaction.
magnesium + oxygen → magnesium oxide.
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Chapter 6
Section 3 Balancing Chemical
Equations
Math Skills, continued
3. Write the equation using formulas for the
elements and compounds in the word equation.
Remember that some gaseous elements, like oxygen, are
molecules, not atoms. Oxygen in air is O2, not O.
Mg + O2 → MgO
4. Balance the equation one element at a time.
The same number of each kind of atom must appear on both
sides. So far, there is one atom of magnesium on each side of
the equation.
But there are two oxygen atoms on the left and only one on
the right.
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Chapter 6
Section 3 Balancing Chemical
Equations
Math Skills, continued
4. Balance the equation one element at a time,
continued
To balance the number of oxygen atoms, you need to
double the amount of magnesium oxide:
Mg + O2 → 2MgO
This equation gives you two magnesium atoms on
the right and only one on the left. So you need to
double the amount of magnesium on the left,
as follows.
2Mg + O2 → 2MgO
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Chapter 6
Section 3 Balancing Chemical
Equations
Math Skills, continued
4. Balance the equation one element at a time,
continued
2Mg + O2 → 2MgO
Now the equation is balanced. It has an equal
number of each type of atom on both sides.
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Chapter 6
Section 3 Balancing Chemical
Equations
Math Skills
Balancing Chemical Equations.
Page 202 Practice 1, 2, 3
Page 204 1-5 (5 is a challenge!)
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Chapter 6
Section 3 Balancing Chemical
Equations
Determining Mole Ratios
• The law of definite proportions states that a
compound always contains the same elements in the
same proportions, regardless of how the compound
is made or how much of the compound is formed.
• Because the law of definite proportions holds true for
all chemical substances in all reactions, mole ratios
can be derived from balanced equations.
• Mole ratio is the relative number of moles of the
substances required to produce a given amount of
product in a chemical reaction.
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Chapter 6
Section 3 Balancing Chemical
Equations
Law of Definite Proportions
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Chapter 6
Section 3 Balancing Chemical
Equations
Determining Mole Ratios, continued
• The mole ratio for any reaction comes from the balanced
chemical equation.
• Example: The equation for the electrolysis of water
shows that the mole ratio for H2O:H2:O2 is 2:2:1.
• 2H2O → 2H2 + [1]O2
• If you know the mole ratios of the substances in a
reaction, you can find the relative masses of the
substances required to react completely.
• Relative masses can be found by multiplying the
molecular mass of each substance by the mole ratio
from the balanced equation.
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Chapter 6
Section 3 Balancing Chemical
Equations
Objectives
• Demonstrate how to balance chemical equations.
• Interpret chemical equations to determine the
relative number of moles of reactants needed and
moles of products formed.
• Explain how the law of definite proportions allows for
predictions about reaction amounts.
• Identify mole ratios in a balanced chemical equation.
• Calculate the relative masses of reactants and
products from a chemical equation.
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Chapter 6
Section 4 Rates of Change
Objectives
• Describe the factors affecting reaction rates.
• Explain the effect a catalyst has on a chemical
reaction.
• Explain chemical equilibrium in terms of equal
forward and reverse reaction rates.
• Apply Le Châtelier’s principle to predict the effect
of changes in concentration, temperature, and
pressure in an equilibrium process.
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Chapter 6
Section 4 Rates of Change
Bellringer
Not all reactions happen at the same speed. Atoms,
ions, and molecules can only interact when they are in
close contact with each other. Below is a sample of zinc
arranged in three different ways.
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Chapter 6
Section 4 Rates of Change
Bellringer, continued
1. In the reaction Zn + 2HCl → ZnCl2 + H2, which sample do you think
would react the fastest? Why?
2. When you want to start a bonfire, why do you use many small sticks as
kindling to start the larger logs?
3. Which do you think will react faster with hydrochloric acid, HCl–atoms of
liquid zinc at its melting point or atoms of solid zinc at its melting point?
(Hint: Which situation allows more contact among the particles?)
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Chapter 6
Section 4 Rates of Change
Factors Affecting Reaction Rates
• For any reaction to occur, the particles of the
reactants must collide with one another. Therefore,
whatever will help particles collide with one another
will speed up the reaction rate.
•
•
•
•
•
Most reactions go faster at higher temperatures.
Greater surface area speeds up reactions.
Concentrated solutions react faster.
Reactions are faster at higher pressure.
Massive, bulky molecules react slower.
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Chapter 6
Section 4 Rates of Change
Factors Affecting Reaction Rate
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Chapter 6
Section 4 Rates of Change
Factors Affecting Reaction Rates, continued
• A catalyst is a substance that changes the rate of a
chemical reaction without being consumed or
changed significantly.
• Catalysts are not reactants or products, because
they are not used up in the reaction.
• Catalysts are often used in industry to make
reactions go faster.
• Catalysts that slow reactions are called
inhibitors.
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Chapter 6
Section 4 Rates of Change
Catalyst
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Chapter 6
Section 4 Rates of Change
Factors Affecting Reaction Rates, continued
• Enzymes are proteins that serve as biological
catalysts.
• An enzyme is very specific, controlling one
reaction or set of similar reactions.
• Most enzymes are fragile, and stop working above
certain temperatures.
• The substrate is the reactant in reactions catalyzed
by enzymes.
• Example: hydrogen peroxide is the substrate for
catalase:
catalase
2H2O2 
 2H2O + O2
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Chapter 6
Section 4 Rates of Change
Inhibitors
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Chapter 6
Section 4 Rates of Change
Enzyme
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Chapter 6
Section 4 Rates of Change
Equilibrium Systems
• Some changes are reversible.
• Example: the physical change represented below can go in
either direction.
increase pressure

 CO2 (gas dissolved in liquid)
CO2 (gas above liquid) 
decrease pressure
• Chemical equilibrium is a state of balance in which
the rate of a forward reaction equals the rate of the
reverse reaction.
• Systems in equilibrium respond to minimize change.
• Example: when the top is removed from a carbonated drink,
the system is no longer at equilibrium, and CO2 leaves as
bubbles.
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Chapter 6
Section 4 Rates of Change
Equilibrium
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Chapter 6
Section 4 Rates of Change
Equilibrium
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Chapter 6
Section 4 Rates of Change
Equilibrium Systems, continued
• Le Châtelier’s principle predicts changes in
equilibrium.
• Le Châtelier’s principle is a general rule that states that if
a change is made to a system in chemical equilibrium,
the equilibrium shifts to oppose the change until a new
equilibrium is reached.
• Le Châtelier’s principle can be used to control
reactions.
• Example: in a reaction that releases energy, if you raise
the temperature, the equilibrium will shift to the left and
make less products.
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Chapter 6
Section 4 Rates of Change
Factors Affecting Equilibrium
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Chapter 6
Section 4 Rates of Change
Le Châtelier’s Principle
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Chapter 6
Section 4 Rates of Change
Concept Mapping
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Chapter 6
Standardized Test Prep
Understanding Concepts
1. Mg(s) + Cl2(g) → MgCl2(s) is an example of what
type of chemical reaction?
A.
B.
C.
D.
synthesis reaction
decomposition reaction
single-displacement reaction
double-displacement reaction
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Chapter 6
Standardized Test Prep
Understanding Concepts
1. Mg(s) + Cl2(g) → MgCl2(s) is an example of what
type of chemical reaction?
A.
B.
C.
D.
synthesis reaction
decomposition reaction
single-displacement reaction
double-displacement reaction
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Chapter 6
Standardized Test Prep
Understanding Concepts
2. Which of the following changes will not increase the
rate of a chemical reaction?
F.
G.
H.
I.
using an enzyme in a reaction
adding an inhibitor to the reaction mixture
increasing the concentration of the reactants
grinding a solid reactant to make a fine powder
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Chapter 6
Standardized Test Prep
Understanding Concepts
2. Which of the following changes will not increase the
rate of a chemical reaction?
F.
G.
H.
I.
using an enzyme in a reaction
adding an inhibitor to the reaction mixture
increasing the concentration of the reactants
grinding a solid reactant to make a fine powder
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Chapter 6
Standardized Test Prep
Understanding Concepts
3. Which of the following is an endothermic chemical
reaction?
A.
B.
C.
D.
fireworks exploding in the sky
photosynthesis in plant cells
respiration in animal cells
wood burning in a fireplace
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Chapter 6
Standardized Test Prep
Understanding Concepts
3. Which of the following is an endothermic chemical
reaction?
A.
B.
C.
D.
fireworks exploding in the sky
photosynthesis in plant cells
respiration in animal cells
wood burning in a fireplace
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Chapter 6
Standardized Test Prep
Understanding Concepts
4. Most chemical reactions proceed faster if the
reactants are heated. How does the added heat
affect reactant atoms or molecules?
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Chapter 6
Standardized Test Prep
Understanding Concepts
4. Most chemical reactions proceed faster if the
reactants are heated. How does the added heat
affect reactant atoms or molecules?
Answer: Addition of heat causes the particles to move
faster and collide more often. The increase in
collisions speeds up the reaction.
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Chapter 6
Standardized Test Prep
Understanding Concepts
5. The reaction of glucose and oxygen to form carbon
dioxide and water produces the same amount of
energy inside living cells as it does by combustion.
Analyze how this reaction can occur at body
temperature in the cells, but not in the open air.
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Chapter 6
Standardized Test Prep
Understanding Concepts
5. The reaction of glucose and oxygen to form carbon
dioxide and water produces the same amount of
energy inside living cells as it does by combustion.
Analyze how this reaction can occur at body
temperature in the cells, but not in the open air.
Answer: Inside living cells, enzymes act as catalysts to
reduce the amount of energy needed to start the
reaction and to allow it to proceed at a lower
temperature.
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Chapter 6
Standardized Test Prep
Reading Skills
Some metals react with water to form new compounds
by displacing hydrogen from water molecules. Alkali
metals are sufficiently reactive that this chemical
reaction happens at room temperature. If a piece of
cesium is placed in water, an explosion occurs as the
hydrogen gas reacts with oxygen in
the air.
continued on next slide
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Chapter 6
Standardized Test Prep
Reading Skills
6. Hydrogen and oxygen gases do not react
spontaneously when they are mixed, unless energy
is added to start the reaction. What is the source of
energy that causes hydrogen to react explosively
when cesium is added to water?
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Chapter 6
Standardized Test Prep
Reading Skills
6. Hydrogen and oxygen gases do not react
spontaneously when they are mixed, unless energy
is added to start the reaction. What is the source of
energy that causes hydrogen to react explosively
when cesium is added to water?
Answer: The reaction of cesium and water is extremely
exothermic. This exothermic reaction provides the
energy to initiate the reaction between hydrogen
and oxygen.
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Chapter 6
Standardized Test Prep
Interpreting Graphics
7. In each of these reactions, the chemical energy increases and
then decreases, during the course of the reaction. What does
the height of the “hill” on each graph represent?
F. energy that must be added to start the reaction
G. energy released as reactant molecules approach one
another
H. the potential energy of the chemical bonds in the molecules
of the reactants
I. The change in total chemical energy between the reactants
and the products
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Copyright © by Holt, Rinehart and Winston. All rights reserved.
Chapter 6
Standardized Test Prep
Interpreting Graphics
7. In each of these reactions, the chemical energy increases and
then decreases, during the course of the reaction. What does
the height of the “hill” on each graph represent?
F. energy that must be added to start the reaction
G. energy released as reactant molecules approach one
another
H. the potential energy of the chemical bonds in the molecules
of the reactants
I. The change in total chemical energy between the reactants
and the products
Chapter menu
Resources
Copyright © by Holt, Rinehart and Winston. All rights reserved.