Chapter 9
Chemical Reactions
1
Vocabulary – Ch. 9.1
•
•
•
•
•
•
•
Chemical reaction
Reactant
Product
Word Equation
Skeleton Equation
Chemical equation
Coefficient
2
Objectives
• Write chemical equations to
describe chemical reactions
• Classify and identify chemical
reactions
• Write ionic equations for
reactions that occur in water
3
Kick Off Activity
• This can be put in your binder. It is
not due until the 4th marking period.
Answer the questions as we cover
the subjects and hold onto it. (-5
points each time you ask for a new
one.)
• This simple reaction that you will do
will summarize most of the things
you will learn over the next few
months.
4
Interpreting Formulas
• Remember: Formulas represents a
formula unit of an ionic compound or the
number of atoms connected in a
molecule.
• How do you represent more than one
formula unit or molecule.
• Add a coefficient (number) before the
formula.
Example: 5 NaCl means 5 units of NaCl
5
Interpreting Formulas
• A formula unit summarizes how
much of each atom is in a compound.
• Multiply the subscript of the element
by the coefficient to get the total
number of each.
Example: 3 NaCl
Means there are three sodium and three
chloride
6
Interpreting Formulas
Practice
Formula of compound
• 2Na2O
Number of each atom
• 4H2SO4
• 5CaSO4•2H2O
7
Writing Chemical
Equations
• In order to understand a chemical
reaction, you have to describe the
changes that take place.
• Part of the description is stating what is
reacting (changing) and what forms.
• Reactants – the substances that
undergo a reaction
• Products – The new substances that
form
8
Writing Chemical
Equations
HC2H3O2 + NaHCO3  CO2 + H2O +
NaC2H3O2
• Make description complete by
indicating the state of the substance.
• Here are the symbols:
(s) = Solid
(l) = Liquid
(g) = Gas
(aq) = Aqueous (substance dissolved in
water)
9
Writing Chemical
Equations Summary
• Here’s a complete chemical formula
description of the reaction we have
witnessed:
HC2H3O2 (aq) + NaHCO3 (s)  CO2 (g) +
NaC2H3O2 (aq) + H2O (l)
Reaction Description
• Aqueous acetic acid and solid sodium
hydrogen carbonate react to form carbon
dioxide gas, aqueous sodium acetate,
and water.
10
Balancing Chemical
Equations
• Balancing chemical equations is the
application of the conservation of
atoms law.
• If the number of atoms on the reactants
(left) side is equal the number of atoms
on the products (right) side, the equation
is balanced.
# Atoms (reactants) = #Atoms (Products)
11
Balancing Chemical
Equations
• Left side of
equation
• HC2H3O2 +
NaHCO3
• Right side of
equation
CO2 + NaC2H3O2 +
H2O
12
Balancing Equations
• Word Equation
Acetic acid (aq) + sodium hydrogen
carbonate (s)  carbon dioxide (g) +
sodium acetate (aq) + water(l)
• Skeleton Equation –describes the
reaction with balanced chemical
formalas
HC2H3O2(aq) + NaHCO3(s)  CO2(g)
+ NaC2H3O2(aq) + H2O(l)
13
Balancing Equations
• Chemical Equation describes the chemical
formulas of the reactants and
products along with their
relative amounts in a
chemical reaction by use of
coefficients
14
Balancing Equations
• Balanced Chemical Equation
HC2H3O2(aq) + NaHCO3(s)  CO2(g)
+ NaC2H3O2(aq) + H2O(l)
Reactants Side
Products Side
5 Hydrogens
5 Hydrogens
3 Carbons
3 Carbons
5 Oxygens
5 Oxygens
1 Sodium
1 Sodium
15
Steps for Balancing
Equations
1. Write the skeleton equation. All
reactants and products have their
correct balanced formulas.
2. By inspection, find an element or
polyatomic group that does not have
equal amounts on both sides.
3. Add coefficients to make the number
of atoms or groups equal on both sides
of equation.
16
Steps for Balancing
Equations
4. Repeat steps 2 and 3 until all atoms
and groups are balanced.
5. Generally, it is better to balance oxygen
and single elements as late as possible.
6. Write coefficients in their lowest
possible whole-number ratios.
17
Balancing Equations
Example
Word Equation:
sodium hydroxide(aq) + carbon dioxide(g) 
sodium carbonate(aq) + water
Skeleton:
NaOH(aq) + CO2(g)  Na2CO3(aq) + H2O(l)
What is the first atom that looks out of balance?
18
Balancing Equations
• Start on this one by balancing sodium on
both sides of the equation.
• How would we do that?
• Skeleton:
NaOH(aq) + CO2(g)  Na2CO3(aq) + H2O(l)
• Put a 2 in front of the NaOH:
2NaOH(aq) + CO2(g)  Na2CO3(aq)+ H2O(l)
19
Balancing Chemical
Equations
Coefficients
• A number written in front of a reactant or
product that states how many units of
something are reacting or being formed.
• Don’t change subscripts in a formula after you
make the skeleton equation!!
• That would change the identity of the
substance!!
• Don’t put coefficients in the middle of a
compound!
20
Balancing Equations
2NaOH(aq) + CO2(g) Na2CO3(aq) + H2O(l)
So what did that coefficient do?
Let’s count the atoms:
Reactants
Products
Na
2
2
O
4
4
H
2
2
C
1
1
• Generally, balance Oxygen and single
element-atoms last.
21
Diatomic
Elements
• Diatomic Molecules –
– Seven non metal elements are found naturally as
molecular elements of two identical atoms
H-H
NΞN O=O
F-F
Cl-Cl
Br-Br
I -I
• They are:
• H2, N2, O2, F2, Cl2, Br2, and I2
– All are gases except Br2 (liquid) and I2 (Solid)
22
Balancing Equations
Example:
Word Equation:
Hydrogen(g) + Chlorine(g) 
Hydrogen chloride(g)
Skeleton Equation:
H2(g) + Cl2(g)  HCl(g)
23
Balancing Equations
Skeleton Equation:
H2(g) + Cl2(g)  HCl(g)
Both H and C are out of balance. Put a 2
in front of HCl to give two of each on both
sides of equation:
H2(g) + Cl2(g)  2HCl(g)
24
Balancing Equations
H2(g) + Cl2(g)  2HCl(g)
Here’s what happened pictorally:
H-H + Cl-Cl  H-Cl and H-Cl
These bonds were broken.
These bonds were formed in making a
new substance.
25
Balancing Equations
Another example:
magnesium(s) + hydrochloric acid(aq) 
magnesium chloride (aq) + hydrogen (g)
Skeleton:
Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)
What do we need to do to balance?
Add a 2 in front of HCl.
26
Balancing Equations
Harder example:
Iron(III) chloride(aq) + sodium hydroxide
(aq)  iron(III) hydroxide(s) + sodium
chloride(aq)
Skeleton & Balanced:
27
Balancing Chemical
Equations - Practice
Word Equations:
1) magnesium chloride (aq) + silver nitrate (aq) 
magnesium nitrate(aq) + silver chloride(s)
2) calcium(s) + water(l)  calcium hydroxide (s)
+ hydrogen(g)
3) lithium phosphate(aq) + magnesium sulfate(aq)
 lithium sulfate(aq) + magnesium phosphate(s)
28
Balancing Chemical
Equations - Practice
1)
magnesium chloride (aq) + silver nitrate (aq)
 magnesium nitrate(aq) + silver chloride(s)
Step 1: Change to a skeleton equation
Step 2: Balance with coefficients
29
Balancing Chemical
Equations - Practice
2)
calcium(s) + water(l)  calcium hydroxide (s)
+ hydrogen(g)
Step 1: Skeleton Equation
Step 2: Balance with coefficients
30
Balancing Chemical
Equations - Practice
3) lithium phosphate(aq) + magnesium
sulfate(aq)  lithium sulfate(aq) +
magnesium phosphate(s)
Step 1: Write Skeleton equation:
Step 2: Balance equation with coefficients
31
Ch. 9.2 - Types of Reactions
•
There are five types of chemical
reactions we will talk about:
1.
2.
3.
4.
5.
•
Synthesis reactions
Combustion reactions
Decomposition reactions
Single replacement reactions
Double replacement reactions
You need to be able to identify the type
of reaction and predict the product(s)
32
Steps to Writing Reactions
•
Some steps for doing reactions
1. Identify the type of reaction
2. Predict the product(s) using the type of
reaction as a model
3. Balance it
Don’t forget about the diatomic elements!
There are seven diatomic elements
They form a “7” on periodic chart and point
toward the 7th (H2)
33
1. Synthesis reactions
Synthesis reactions occur when two
substances (generally elements) combine and
form a compound. (Sometimes these are called
combination, formation or addition reactions.)
reactant + reactant  1 product
• Basically: A + B  AB
• Example: 2H2 + O2  2H2O
• Example: 2Mg + O2  2MgO
• Example: CaO(s) + H2O(l)  Ca(OH)2(s)
34
Synthesis Reactions
• Here is another example of a
synthesis reaction
35
Practice
Predict the products. Write and balance the
following synthesis reactions.
1. sodium(s) + chlorine(g) 
2. magnesium(s) + Iodine(s) 
3. aluminum(s) + fluorine(g) 
36
2. Combustion Reactions
• Combustion reactions occur
when oxygen combines with
another substance and releases
energy as heat and light.
• Example:
Mg(s) + O2(g)  MgO(s)
37
Combustion Reactions
Combustion is also called
burning!!! In order to burn
something you need the
3 things in the “fire
triangle”:
1) A Fuel (hydrocarbon)
2) Oxygen to burn it with
3) Something to ignite
the reaction (spark)
38
Combustion Reactions
• In general:
CxHy + O2  CO2 + H2O
• Hydrocarbon combustion
ALWAYS produces carbon
dioxide and water. (Oxide of
carbon and oxide of hydrogen)
• Combustion is used to heat
homes and run automobiles
(octane, as in gasoline, is
C8H18)
39
Combustion
• Example
•
C5H12 + 8 O2  5 CO2 + 6 H2O
• Write the products and balance the
following combustion reaction:
•
C9H20 +
O2 
40
Classifying Chemical
Reactions - Combustion
• Where can we find examples of a
combustion reaction?
•
•
•
•
•
– Combust CH4 to heat homes
– Combust gasoline for transportation
Another example:
Ethanol (l) + oxygen(g)  carbon dioxide(g) +
water(g)
C2H5OH(l) + O2 (g) → CO2(g) + H2O (g)
C2H5OH (l) + 3O2(g) → 2CO2(g) + 3H2O (g)
What evidence do you have that reaction
occurred?
41
3. Decomposition
Reactions
• Decomposition reactions occur when a
compound breaks up into the elements or
in a few to simpler compounds
• 1 Reactant  Product + Product
• In general: AB  A + B
• Example: 2 H2O  2H2 + O2
• Example: 2 HgO  2Hg + O2
42
Decomposition
Reactions
• Another view of a decomposition
reaction:
43
Decomposition
Exceptions
• Carbonates and chlorates are special case
decomposition reactions that do not go to
the elements.
• Carbonates (CO32-) decompose to carbon
dioxide and a metal oxide
• Example: CaCO3  CO2 + CaO
• Chlorates (ClO3-) decompose to oxygen gas
and a metal chloride
• Example: 2 Al(ClO3)3  2 AlCl3 + 9 O2
• There are other special cases, but we will not
explore those in Chemistry I
44
Practice
• Predict the products. Then, write
and balance the following
decomposition reaction equations:
• Solid Lead (IV) oxide decomposes
PbO2(s) 
• Aluminum nitride decomposes
AlN(s) 
45
Classifying Chemical
Reactions - Decomposition
• What happens when things
decompose?
– They break down into simpler structures
– Take for example, NaHCO3
– Can you name this compound?
– NaHCO3(s) → Na2CO3(s) + CO2(g) + H2O(g)
– 2NaHCO3(s)→ Na2CO3(s) + CO2(g)+ H2O(g)
46
Practice
Identify the type of reaction (synthesis,
combustion or decomposition) and
balance the equation:
1. __Al2O3(s)  __Al(s) + __O2(g)
2. __C3H8(g)+ __O2(g)__CO2(g)+__H2O(g)
3. __Co(s)+ __S(s)  Co2S3(s)
4. __BaCO3(s) __BaO(s) + __CO2(g)
47
4. Single Replacement
Reactions
Single Replacement Reactions occur when
one element replaces another in a
compound.
A metal can replace a metal ion (+) OR
a nonmetal can replace a nonmetal ion (-).
element + compound product + product
A + BC  AC + B (if A is a metal) OR
A + BC  BA + C (if A is a nonmetal)
(remember the cation always
goes first!)
48
Single Replacement
Reactions
• Another view:
49
Single Replacement
Reactions - Example
• Write and balance the following single
replacement reaction equation:
• Zinc(s) + hydrochloric acid(aq)  zinc
chloride(s) + hydrogen(g).
Zn(s) + 2HCl(aq)  ZnCl2(s) + H2(g)
Note: Zinc replaces the hydrogen ion in
the reaction
50
Single Replacement
Reactions
• Reaction:
51
Single Replacement
Reactions (Redox)
• In reality, ‘Single replacement
reaction’ is a simplified term for
a REDOX reaction
• Stands for:
– Oxidation: Loss of electrons
Example: Zn(s)  Zn2+(aq) + 2e-
– Reduction: Gain of electrons
Example: 2H+(aq)+ 2e-  H2(g)
52
Single Replacement
(Redox) Reactions
• Oxidation Number is the charge on
an atom (zero) or ion.
• So in our previous example:
Zn(s)  Zn2+(aq) + 2e• The Oxidation number of Zn(s) metal is
zero.
• The Oxidation number of the Zn2+ ion is
…+2.
• Oxidation number increases (more
positive) for an oxidized substance.
53
Single Replacement
Reactions (Redox)- Examples
• Sodium chloride(s) + fluorine(g)  sodium
fluoride(s) + chlorine(g).
2 NaCl(s) + F2(g)  2 NaF(s) + Cl2(g)
Note that fluorine replaces chlorine in the compound
• Aluminum(s) + copper(II) nitrate(aq) 
aluminum nitrate(aq) + copper(s)
2Al(s)+ 3Cu(NO3)2(aq)
2Al(NO3)3(aq) + 3Cu(s)
54
Examples of Activity
• AgNO3 (ag) + Cu(s) 
•
Cu(NO3)2(aq) + Ag(s)
• Copper metal is more active than silver
ion so it replaces silver ion in solution.
• Ag(s) + CuSO4(aq)  No Reaction
• Why? Because copper ion is more active
than silver metal and will remain as an
ion.
• Note that activity for metals means it is
more likely to be oxidized.
55
Practice
• Predict the products, if any, note
oxidation #’s, and balance the reaction:
1. __Fe(s) + __CuSO4(aq) 
__FeSO4(aq) + __Cu(s)
2. __Br2(l) + __MgCl2(aq)  No Reaction
3. __Fe(s) + __Na3PO4(aq)  No Reaction
4. __Al(s) + _3CuCl(aq) 
__AlCl3(aq) + 3_Cu(s)
56
Chapter 7 Review
Name
• Sodium chloride
Formula
• Potassium
carbonate
• Molybdenum(VI)
chloride
• Iron(III) oxalate
57
5. Double Replacement
Reactions
Double Replacement Reactions occur
when a metal ion replaces a metal ion in a
compound and a nonmetal ion replaces a
nonmetal ion in a compound.
Compound + compound  product +
product
• AB + CD  AD + CB
58
Classifying Chemical
Reactions – Double
Replacement
• Example:
• CaCl2(aq) + K2CO3(aq) →
CaCO3(s)+ 2KCl(aq)
59
Double Replacement
Reactions
• Think about it like “foil”ing in algebra, first and
last ions go together + inside ions go together
• Example:
AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)
• Another example:
K2SO4(aq) + Ba(NO3)2(aq) 2 KNO3(aq) + BaSO4(s
60
Practice
• Balance the following equations.
1. HCl(aq) + AgNO3(aq) 
AgCl(s) + HNO3(aq)
2. 3CaCl2(aq) + 2Na3PO4(aq) 
Ca3(PO4)2(s) + 6NaCl(aq)
3. FeCl3(aq) + 3NaOH(aq) 
Fe(OH)3(s) + 3NaCl(aq)
61
Harder Practice
•
Predict products, then balance the
following equations.
1. Pb(NO3)2 + BaCl2 
PbCl2 + Ba(NO3)2
2. H2SO4 + 2NaOH 
2HOH + Na2SO4
3. 2KOH + CuSO4 
K2SO4 + Cu(OH)2
62
Reactions in Aqueous
Solutions – Ch. 9.3
• Solute is dissolved in water – ions
are separated from each other
(eg CaCl2 is a solute)
• Solvent is what it is dissolved in (eg
water is the solvent)
• These form an aqueous solution
(aq)
• The solute particles dissociate, they
separate in the solution.
63
Complete Ionic Equations
• What really happened in our famous
experiment?
• Solutes CaCl2 and K2CO3 are soluble in
water, so
• The reactants CaCl2(aq) and
K2CO3(aq) dissociated to make ions
Ca+2(aq) and 2Cl-(aq) and 2 K+(aq) and
CO3-2(aq)
• Then what happened?
64
Complete Ionic Equations
• Ca+2 (aq) and CO3-2 (aq) got together
to form a solid CaCO3(s) leaving
2K+(aq) and 2Cl-(aq) still dissolved in
solution.
• How do we know CaCO3 forms a
solid? See solubility table.
65
Describing Ionic Equation
for CaCO3 Reaction
• Complete ionic equation shows all
particles as they exist in solution.
Ca+2(aq) + 2Cl- (aq) + 2K+ (aq) + CO3-2 (aq)
 CaCO3(s) + 2 K+(aq) + 2Cl-(aq)
• Note that Cl- and K+ don’t change! They
are spectator ions.
• Net ionic equation shows only the
particles that take part in the reaction.
Ca+2(aq) + CO3-2(aq) → CaCO3(s)
66
Net Ionic Equations
• Cancel out ions that appear on BOTH sides of
the reaction in the complete ionic equation to
get the net ionic equation.
Complete Ionic Equation:
2 K+(aq) + CrO4 -2(aq) + Pb+2(aq) + 2 NO3-(aq)
 PbCrO4(s) + 2 K+(aq) + 2 NO3-(aq)
Net Ionic Equation:
CrO4 -2(aq) + Pb+2(aq)  PbCrO4(s)
67
Net Ionic Equations Practice
• From the full chemical equation write the
complete ionic and net ionic equations for this
reaction.
Chemical:
2AgNO3(aq) + K2Cr2O7(aq) Ag2Cr2O7(s) + 2KNO3(aq)
Complete Ionic:
Net Ionic:
68
More Net Ionic Equations
Practice
Given this Full Chemical Equation:
K2CrO4 + Pb(NO3)2  PbCrO4(s) + 2 KNO3
Soluble
Soluble
Insoluble
Soluble
Write the Complete Ionic Equation:
And the Net Ionic Equation:
69
Reactions in Aqueous
Solutions
• If all solutes are soluble and don’t react to
produce a solid or a gas then no reaction
occurs. Example are solutes KCl and
Zn(NO3)2:
K+ (aq) + Cl-(aq) + Zn+2(aq) + 2NO3-(aq)
K+(aq) + Cl-(aq) + Zn+2(aq) + 2NO3-(aq)
Which really means:
K+ (aq) + Cl-(aq) + Zn+2(aq) + 2NO3-(aq)  NR
70
Water formed in Reaction
• When a solid is one of the products, it is
called a precipitate. We’ve had numerous
examples CaCO3, PbCrO4 and AgCl
• Can also make water if have H+ and OH- in
solution (acid/base)
– H+(aq) + NO3-(aq) + Na+(aq) + OH-(aq)
 Na+(aq) + NO3-(aq) + H20(l)
• Or H+(aq) + OH-(aq)  H2O(l)
71
Gases formed in Ionic
Reaction
• Sulfuric acid (H2SO4) reacts with 2NaCN:
• Complete Ionic Equation:
• 2H+(aq) + SO4-2(aq) + 2Na+(aq) + 2CN- (aq)
 2HCN(g)+ SO4-2(aq) + 2Na+(aq)
• Net ionic equation:
– H+(aq) + CN-(aq)  HCN(g)
72