Chapter 8
Reactions and Aqueous Solutions
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Objectives
• To learn about some of the factors that cause reactions to
occur.
• To learn to identify the solid that forms in a precipitation
reaction.
• To learn to describe reactions in solutions by writing
molecular, complete ionic, and net ionic equations.
• To learn the key characteristics of the reactions between
strong acids and strong bases.
• To learn the general characteristics of a reaction between a
metal and a nonmetal.
• To understand electron transfer as a driving force for a
chemical reaction.
• To learn various classification schemes for reactions.
• To consider additional classes of chemical reactions.
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8.1 Predicting Whether a Reaction will
Occur
• Why do chemical reactions occur? What causes
reactants to “want” to form products?
• Driving forces pull reactants towards products.
These are changes that tend to make reactions go
in the direction of the arrow.
• Most common driving forces are:
– Formation of solid, water, or gas, and transfer of
electrons.
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8.2 Reactions in Which a Solid Forms
• Reactions which solids form are known as precipitation
reactions.
• The solid that forms during the reaction is known as a
precipitate.
• If you think back to your 8th grade science class, you will
remember that precipitates are a sign of a chemical
change. Other signs include:
– Color change
– Odor change
– Gas, heat, light, sound production
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What happens when an ionic compound
dissolves in water?
• In virtually every case when a solid containing
ions dissolves in water, the ions separate and
move around independently.
• This is process is known as dissociation.
• When looking at a chemical equation, these
are the aqueous reactants and products.
Example: Ba(NO3)2(aq) dissociates and becomes
Ba2+ and 2 NO3- ions when dissolved in water.
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How can we tell?
• When we try and conduct electricity in pure
water, what happens?
• What if we dissolve ions in the water?
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Strong Electrolytes
• Ions that completely dissociate in the water have
positive and negative charges which allows
electrons to flow through the water and conduct
electricity.
• Strong electrolytes are materials that, when
dissolved in water, dissociate completely giving a
solution that conducts electricity very well.
• Conversely, when sugar is dissolved in water, it
does not dissociate and electricity is not
conducted. Sugar is not an ionic compound.
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How to Predict the Products
• If we have 2 compounds and dissolve them in
water, how will we know what the products
will be?
Example:
Ba(NO3)2(aq) + K2CrO4(aq)  ?????
• First let’s determine the ions that will be
present in a solution:
Ba2+ + 2NO3- + 2K+ + CrO428
Consider all the Possible Products
• If we set up a table we can figure out all of the
possible products, since we can only have cations and
anions together.
NO3-
CrO42-
K+
Ba2+
• One of these compounds is going to form a yellow
precipitate, but how we decide which one it will be?
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Solubility Rules
• There are 2 types of solids:
• Soluble – solid that readily dissolves in water
Example: salt, NaCl
• Insoluble – a solid where a tiny amount dissolves
in water that is undetectable to the naked eye,
thus we can still see a solid in the liquid.
Example: any visible precipitate
To help determine what will be soluble and
insoluble we have a set of solubility rules.
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Back to Our Product
• So, if we look at our
reaction and possible
products we can use
our solubility rules to
determine the
precipitate.
K+
Ba2+
• Most nitrates are
soluble
• Most salts of K+ are
soluble
• Which leaves us with
BaCrO4CrO 2-
NO3KNO3
Ba(NO3)2
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K2CrO4
BaCrO4
The overall reaction would look like this:
K2CrO4(aq) + Ba(NO3)2(aq)  BaCrO4(s) + 2KNO3(aq)12
Practice
• Predict which of the following would be soluble in
water:
–
–
–
–
Potassium nitrate
Zinc hydroxide
Calcium hydroxide
Ammonium chloride
• Use the solubility rules to predict the precipitate
product between the following and write the balanced
equation for the reaction:
– KCl(aq) + AgNO3(aq)
– KOH(aq) + Fe(NO3)3(aq)
– Na2SO4(aq) + Pb(NO3)2(aq)
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8.3 Describing Reaction in Aqueous
Solutions
• There are 3 ways we can describe a reaction in
water such as the one below:
K2CrO4(aq) + Ba(NO3)2(aq)  BaCrO4(s) + 2KNO3(aq)
• This equation is called the molecular equation.
• The molecular equation shows the overall reaction,
but not necessarily the actual reactants and products
in solution.
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Complete Ionic Equations
• If we want to look at what is occurring in the
solution we write the complete ionic equation.
• We write all aqueous compounds as the ions they
would appear as in a solution.
K2CrO4(aq) + Ba(NO3)2(aq)  BaCrO4(s) + 2KNO3(aq)
Becomes:
2K+(aq) + CrO42-(aq) + Ba2+(aq) +2NO3-(aq) 
BaCrO4(s) + 2K+(aq) + 2NO3-(aq)
• In a complete ionic equation, all reactants and
products that are strong electrolytes as ions. All
reactants and products are included.
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Net Ionic Equation
• The complete ionic equation reveals that only
some of the ions participate in the reaction.
2K+(aq) + CrO42-(aq) + Ba2+(aq) +2NO3-(aq) 
BaCrO4(s) + 2K+(aq) + 2NO3-(aq)
• The ions that did not change on either side of
the equation are known as spectator ions.
• We can write the equation without these ions,
as they do not directly affect the reaction.
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Net Ionic Equations
• If we remove the spectator ions from the complete
ionic equation, we get the net ionic equation.
Ba2+(aq) + CrO42-(aq)  BaCrO4(s)
• The net ionic equation includes only those
components that are directly involved in the
reaction, spectator ions are not included.
• It is important to note that when doing this
process only aqueous compounds are changed to
ions. Solids, liquids and gases do not become
ions.
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Practice
• For each write the molecular, complete ionic, and
net ionic equation.
– Aqueous sodium chloride is added to aqueous silver
nitrate to form solid silver chloride and aqueous
sodium nitrate.
– Aqueous potassium hydroxide is mixed with aqueous
iron(III) nitrate to form solid iron(III) hydroxide and
aqueous potassium nitrate.
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8.4 Reactions that Form Water: Acids
and Bases
• Acids have been known to exist for hundreds
of years, characterized by their sour taste the
word acid is derived from the Latin word
acidus, meaning sour.
• Bases are commonly known for their bitter
taste, and slippery feel. Bases are commonly
used in commercial drain cleaners, such as
Draino.
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History
• Acids have been around for hundreds of years, but until
the 1800’s relatively little was known about them.
• Acids were originally found in minerals and treatment of
minerals. These include sulfuric acid and nitric acid.
• In the early 1880s a Swedish graduate student named
Svante Arrhenius discovered the nature of how acids
function. His ideas were scorned until the discovery
that atoms contain charged particles in the late 1890s.
Arrhenius was eventually given the Nobel Prize for his
work in 1903.
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So what did Arrhenius do?
• Arrhenius was studying the reason why some
solutions conduct electric current and others
were not.
• During his studies Arrhenius discovered that
when HCl, H2SO4, and HNO3 were dissovled in
water they dissociated and behaved as strong
electrolytes.
• He suggested that these reactions must look like
the following in water:
HCl  H+(aq) + Cl-(aq)
HNO3  H+(aq) + NO3-(aq)
H2SO4  H+(aq) + HSO4-(aq)
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Strong Acids
• From these results Arrhenius defined an acid to
be the following:
An acid is a substance that produces H+ ions (protons)
when it is dissolved in water.
When HCl, HNO3, and H2SO4 are dissolved in water
100% dissociates to give ions.
Of course these are known as strong electrolytes,
and due to the H+ ions produced we call these
strong acids.
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Strong Bases
• If there are strong acids then there must also
be strong bases.
• Arrhenius defined a base as a substance that
produces hydroxide ions (OH-) in water.
• A strong base is a base that dissociates 100%
when dissolved in water, just like strong acids.
• Example: NaOH(s)  Na+(aq) + OH-(aq)
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Indicating Acids and Bases
• We use indicators to identify acids and bases in a
solution.
• There are many different indicators that tell you
how many H+ ions are present in a solution.
• Example:
• Phenolphthalein is an indicator that indicates when bases are
present.
• This means that there are more OH- ions present in the
solution than H+ ions.
• Phenolphthalein turns pink in the presence of bases.
• Many other indicators exist: pink and blue indicator
paper and cabbage juice are 2 you may have used
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Acid-Base Reactions
• Write the molecular equation, complete ionic,
and net ionic equations for the reaction of
aqueous HCl with aqueous NaOH.
• HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq)
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Acid-Base Reactions
• There are some defining things about acid base
reactions.
• HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
• There will always be H2O(l) in the product.
• There will always be a salt (NaCl)in the product.
• H+(aq)+Cl-(aq)+Na+(aq)+OH-(aq)Na+(aq)+Cl-(aq)+
H2O(l)
• Remove the spectator ions.
• H+(aq) + OH-(aq)  H2O(l) This is the net ionic
equation for all acid base reactions.
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Practice
• Write the molecular, complete ionic and net
ionic equations for the following reaction:
HNO3(aq) + KOH(aq) 
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8.5 Reactions of Metals and Nonmetals
(Oxidation-Reduction or Redox Reactions)
• So thus far we have identified 2 types of reactions:
– Reactions that form a solid
– Reactions that form water
• This section deals with reactions involving a metal
and a non-metal. What type of bond occurs
between a metal and a non-metal?
Ionic bonds – atoms are giving and taking electrons.
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Transfer of Electrons
• So let us look at an example reaction:
2Na(s) + Cl2(g)  2NaCl(s)
• How many electrons does each atom in the
reactants have?
Na has 11 electrons
Cl has 17 electrons
In the ionic compound NaCl, Na has a +1 charge
and Cl has a -1 charge. How did this happen?
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Electron Transfer
• In this reaction the sodium atom gives one
electron to the chlorine atom. This results in the
atoms becoming ions:
Na loses 1 electron, now has 10 electrons and is Na+
Cl gains 1 electron, now has 18 electrons and is Cl-
• Any reaction that involves the transfer of electrons
is known as an oxidation-reduction reaction. Or a
redox reaction.
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Characteristics of Oxidation-Reduction
Reactions
1.
A metal-nonmetal reaction can always be assumed to be an
oxidation-reduction reaction, which involves electron transfer.
When a metal reacts with a nonmetal, an ionic compound is
formed.
The ions are formed when the metal transfers one or more
electrons to the nonmetal.
2. Two non-metals can also undergo an oxidation-reduction reaction.
At this point we can recognize these cases only by looking for O2 as
a reactant or product.
When 2 non-metals react, the compound formed is not ionic.
Example: CH4(s) + 2O2(g)  CO2(g) + 2H2O(g)
2SO2(g) + O2(g)  2SO3(g)
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Practice
• Look at the reactions below and determine
how the electrons are gained and lost.
– 2Na(s) + Br2(l)  2NaBr(s)
– 2Ca(s) + O2(g)  2CaO(s)
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8.6 Ways to Classify Reactions
• Thus far we have considered the following
reactions:
– Formation of a solid = Precipitation Reaction
– Formation of water = Acid/Base Reaction
– Transfer of electrons = Oxidation-Reduction
• However, there is another set of ways that we
sometimes classify reactions.
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Double Displacement Reaction
• Precipitation and acid base reactions can be classified as
double displacement reactions:
• The general form for double displacement reaction is:
AB + CD  AD + BC
• Any reaction which the ions exchange places is called a
double displacement reaction.
Ex: K2CrO4(aq)+Ba(NO3)2(aq)BaCrO4(s)+2KNO3(aq)
Ex: NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(l)
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Single Replacement Reactions
• Like a double displacement reaction, but only
the anions (negative ions) exchange places.
• General form: A + BC  B + AC
• Oxidation reduction reactions can often be in
this form:
Ex: Zn(s) + 2HCl(aq)  H2(g) + ZnCl2(aq)
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8.7 Other Ways to Classify Reactions
• So far we have many types of chemical reactions:
– Precipitation Reactions
– Acid Base Reactions
– Oxidation-Reduction Reactions
– Double Displacement Reactions
– Single Displacement Reactions
• Many times reactions can be classified more than
one way. This section we will discuss a few more
ways reactions can be classified.
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Combustion Reaction
• Combustion reactions are reactions that involve
oxygen and produce energy..
Example: CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
We have discussed this reaction before, what type of
reaction was it?
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Combustion as Oxidation Reduction
• Combustion reactions are also oxidationreduction reactions.
• The oxidation term in the name refers to the
oxygen present in combustion reactions.
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Reaction Map
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Synthesis (Combination) Reactions
• Synthesis or combination reactions are reactions
which a given compound is formed from simpler
materials.
• Examples:
2H2(g) + O2(g)  2H2O(l)
C(s) + O2(g)  CO2(g)
Mg(s) + F2(g)  MgF2(s)
6CO2 + 6H2O  C6H12O6 + 6O2
So you see that each of these forms one major product,
but also they are all another type of reaction that we
have previously discussed.
Oxidation-Reduction Reactions
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Decomposition Reactions
• Decomposition reactions are reactions which a
compound is broken down into simpler compounds or
elements
• Examples:
– 2H2O(l)  2H2(g) + O2(g) (electrolysis)
– 2HgO(s)  2Hg(l) + O2(g) (heat)
– 2NaCl(s)  2Na(l) + Cl2(g) (electric current)
• Each of these reactions break down one large
compound into smaller pieces.
• Each of these can also be classified as oxidationreduction reactions as well.
• Electron transfer is occurring, just in the reverse order.
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Overall Map of Chemical Reaction
Types
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Summary of Reaction Characteristics
• Precipitation – solid precipitate forms, all others are aqueous
• Acid-Base – liquid water and a salt are always in the products
• Double Displacement – AB + CD  AD + BC, can also be precipitation
or acid-base
• Oxidation-Reduction – electrons are transferred between 2 or more
atoms, O2 is often involved
• Single Displacement – A + BC  B + AC, can also be oxidation
reduction
• Combustion – oxygen involved in producing energy, also an oxidationreduction reaction
• Synthesis – one compound is formed from many simpler materials,
also an oxidation-reduction reaction
• Decomposition – one compound is broken down into simpler
compounds or elements, also an oxidation-reduction reaction
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Practice
• Classify each of these reactions in as many ways as
possible:
•
•
•
•
•
•
•
•
2K(s) + Cl2(g)  2KCl(s)
Fe2O3(s) + 2Al(s)  Al2O3(s) + 2Fe(s)
2Mg(s) + O2(g)  2MgO(s)
HNO3(aq) + NaOH(aq)  H2O(l) + NaNO3(aq)
KBr(aq) + AgNO3(aq)  AgBr(s) + KNO3(aq)
PbO2(s)  Pb(s) + O2(g)
4NH3(g) + 5O2(g)  4NO(g) + 6H2O(g)
BaCl2(ag) + Na2SO4(aq)  BaSO4(s) + 2NaCl(aq)
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