Unit 7: Chemical Reactions
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
endothermic and exothermic reactions.
• Identify situations involving chemical energy.
What are chemical reactions?
• As we know, there are two types of changes that
matter can undergo: physical and chemical.
• Recall that a physical change may change the
appearance of a substance, but does not change
its composition.
• However, when a substance undergoes a
chemical change, a new substance with new
properties is formed.
• A chemical reaction is a chemical change.
Signs of chemical reactions
• Often, there are visible signs that a chemical
reaction has taken place.
• Things to look for include:
– Production of gas bubbles (such as CO2).
– The formation of a solid precipitate (heavier
substance that sinks to the bottom of the
substance)
– The production of light or heat
– Change of color as a result of heating
• The parts of a chemical reaction can be split into two
categories: the products and the reactants
• The reactants are the substances that a reaction
begins with
• The products are the substances that are formed in
the chemical reaction
• The products and reactants contain the same types
and quantity of atoms, but they have different
arrangements
• Remember: you can’t create or destroy matter, so the
atoms you end up with have to come from the atoms
you start with
Energy and Reactions
• As you may remember from the previous units,
compounds are different from mixtures in
several ways. One of these is that you can’t just
pull the pieces of a compound apart like you can
with a mixture.
• In fact, the only way to separate the parts of a
compound is to add energy to break the bonds
between the atoms.
• This energy can be heat, light, sound, or
electricity.
• Likewise, when a bond is formed, energy is
released. Again, this energy can be in the form
of light, heat, sound, or electricity.
• If a reaction releases energy, it is said to be
exothermic.
• If a reaction takes in energy, it is endothermic.
EXOTHERMIC
ENDOTHERMIC
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.
Types of reactions
• There are 5 types of chemical reactions:
– Synthesis reactions
– Decomposition reactions
– Combustion reactions
– Single-displacement reactions
– Double-displacement reactions
Synthesis reactions
• In a synthesis reaction, many small molecules
join to form fewer larger molecules
(polymers).
• A + B  AB is the general form
• 2Na + Cl2  2NaCl is a common example
Decomposition reactions
• A reaction in which large substances are
broken apart into smaller atoms or molecules
• AB  A + B is the general form
• 2H2O  2H2 + O2 is a common example
Combustion reactions
• Require oxygen as a reactant, so at least one
product of the reaction contains oxygen
• This usually involves organic compounds, and
heat is often released
• There is no general form for this type of
reaction, except that O2 is always a reactant
• 2CH4 + 4O2  2CO2 + 4H2O is a common
example, where methane combines with
oxygen in the air to form carbon dioxide and
water
Single-displacement reactions
• When the atoms of one element appear to
move into a compound, and atoms of the
other element appear to move out, it is a
single-displacement
• AX + B  BX + A is the general form
• 3CuCl2 + 2Al  2AlCl3 + 3Cu is a common
example
Double-displacement reactions
• Here, ions appear to be exchanged between
compounds
• AX + BY  AY + BX is the general form
• Pb(NO3)2 + K2CrO4  PbCrO4 + 2KNO3 is a
common example
Name That Reaction!
A)
B)
C)
D)
S8 + 8O2  8SO2 + heat
6CO2 + 6H2O  C6H12O6 + 6O2
2NaHCO3  Na2CO3 + H2O + CO2
Zn + 2HCl  ZnCl2 + H2
A)
B)
C)
D)
synthesis
double-displacement
decomposition
single-displacement
Section 3: Balancing Chemical
Equations Objectives
• Demonstrate how to balance chemical
equations.
• Interpret chemical equations to determine the
relative number of moles of reactant needed
and moles of products formed.
• Explain how the law of definite proportions
allows for predictions about reaction
amounts.
Chemical Equations
• A chemical equation is used to describe what
happens during a chemical reaction.
• It makes use of the element symbols and
compound formulas that we’ve already learned.
• The reactants are always on the left-hand side
and the products are always on the right-hand
side. The arrow () means “yields”.
Reactants  Products
Conservation of Mass
• Remember that matter cannot be created or
destroyed, according to the law of
conservation of mass.
• That means that we can’t make more atoms
than we start with, and we can’t lost atoms
along the way.
• So, the number of atoms in the reactants
must equal the number of atoms in the
products.
• Balancing equations helps us to do this.
How to balance chemical
equations
• A few things to keep in mind when trying to balance
equations:
– You cannot change the formula of the compound or molecule
– You cannot change the subscripts in the compound or molecule
– You can use coefficients (numbers placed in front of a
compound or molecule to tell how many there are) to balance
– A coefficient multiplies everything in that compound or
molecule by the given number
– Anything in parentheses acts like a single atom (remember
polyatomic ions?) and if you multiply one part of it, you must
multiply all of it
• Start with:
CH4 + O2  CO2 + H2O
Notice that on the reactant side, there is one C,
two O, and four H
On the product side, there is one C, three O (two
of them in the first compound and one in the
second), and two H.
Therefore, this equation is not balanced.
• To balance this equation, we must use
coefficients. Let’s start with the H atoms.
CH4 + O2  CO2 + 2H2O
The reactant side now has one C, four H and two O.
The product side now has one C, four H, and three
O.
We’re almost balanced!
• Now, to balance the O atoms, we must do one
more thing:
CH4 + 2O2  CO2 + 2H2O
This gives us one C on each side of the arrow,
four H on each side, and four O on each side.
Notice that the 2 in front of H2O multiplies
both the H and the O.
Our equation is balanced!
• A balanced equation demonstrates the law of
conservation of mass because the combined masses
of all the atoms on the reactant side now equals the
masses of all the atoms on the product side
• A balanced equation also tells us how many molecules
of a substance are required for a reaction to take
place. In the example from before, methane will only
combust if two molecules of oxygen gas are present,
and will always produce (or yield) one molecule of
carbon dioxide and two molecules of water. That
reaction will always happen in the same proportions.
Let’s try another!
• Mg + O2  MgO
Right now, there is one atom of Mg and two
atoms of O on the reactant side, and only one
atom of each on the product side.
Therefore, we must balance!
• The next step is to decide where we need to
place coefficients. Remember that a
coefficient goes in front of a compound or
molecule. You cannot separate a compound,
nor can you change the formula or subscripts.
• __Mg + __O2  __MgO
• By placing a 2 in front of the MgO, we now
get:
__Mg + __O2  2MgO
This means we now have one Mg and two O
to start with, and two Mg and two O to end
with.
Not quite finished!
• To balance out the O in our final step, we must
use a coefficient of __. This will give us:
2Mg + __O2  2MgO
• Now, let’s try a little trickier one.
Pb(NO3)2 + KI  KNO3 + PbI2
What’s different about this equation?
Let’s count atoms. On the reactant side, we
have one Pb, two N, six O, one K, and one I.
On the product side, we have one K, one N,
three O, one Pb, and two I.
Definitely needs some work!
• To continue, we need to pick an element and
start balancing. Let’s start with our I atoms.
__Pb(NO3)2 + 2KI  __KNO3 + __PbI2
By placing a 2 in front of the KI, we have
doubled our number of K and I. That means
we now have one Pb, two N, six O, two K, and
two I as reactants, and one Pb, one N, three
O, one K, and two I as products.
Almost there.
• Now let’s balance our K atoms.
__Pb(NO3)2 + 2KI  2KNO3 + __PbI2
By placing a 2 in front of our KNO3 atom, we
now have one Pb, two N, six O, two K, and two
I as reactants, and one Pb, two N, six O, two K,
and two I as products.
Success!
The Law of Definite Proportions
• The law of definite proportions says 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.
This means that water will always have twice as
many H as O, whether it’s made in a synthesis
reaction, decomposition, combustion, or
displacement, and regardless of what it’s reacting
with.
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 LeChatelier’s principle to predict the
effect of changes in concentration,
temperature, and pressure in an equilibrium
process.
• The conditions under which most chemical
reactions take place are very specific. For
many of them, they can only occur if the
temperature is just right, or if there is the right
amount of other substances present, or if the
wind is blowing in the right direction.
• Ok, maybe not that last one, but the others
are true.
• Other things that can affect reaction rate are
surface area, concentration of the substance,
pressure, size of the molecules and the presence
of catalysts.
• In general, higher temperature, surface area,
concentration, and pressure make a reaction go
faster.
• Heavier molecules usually react more slowly, and
catalysts (enzymes or chemicals that can speed
up a reaction by reducing the amount of energy
needed) make reactions occur more quickly.
Equilibrium
• Not all reactions are terminal
• That means that sometimes, a reaction is not
finished when the products are made.
Sometimes, the reaction will start over
backwards.
• This is called an equilibrium system.
• The reaction can go in either direction,
depending on conditions such as temperature,
amount of pressure, and amount of the
substances involved.
LeChatelier’s Principle
• LeChatelier’s Principle says 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.
This principle can be used to control reactions
because those conditions can be manipulated
(such as increasing temperature or decreasing
pressure) to cause the reaction to go in the
direction you desire.