Balancing Equations
– Chemical reactions occur when bonds (between
the electrons of atoms) are formed or broken
– Chemical reactions involve

changes in the chemical composition of matter
– the making of new materials with new properties

energy changes:
– Bond breaking absorbs Energy (endothermic process)
– Bond making releases Energy (exothermic process)
– Symbols represent elements
– Formulas describe compounds
– Chemical equations describe a chemical
reaction
Chemical Equations
A chemical equation is written as an
expression similar to a mathematic
equation that can be compared to a
recipe that a chemist follows in order
to produce desired results.
Chemical Equations
Their Job: Depict the kind of reactants
and products and their relative
amounts in a reaction.
4 Al (s) + 3 O2 (g) ---> 2 Al2O3 (s)
The numbers in the front are called
stoichiometric coefficients
The letters (s), (g), and (l) are the
physical states of compounds.
Chemical Equations

Because of the principle of the
conservation of matter
(matter can not be created
or destroyed) an equation
must be balanced.

It must have the same number of
atoms of the same kind on both sides.
 Law of Conservation of Energy MUST
ALSO BE FOLLOWED!
 Energy changes are written in (endo-/
exothermic reactions)
Lavoisier, 1788
Chemical Equations

All chemical equations have reactants and
products.
 We express a chemical equation as follows:
Reactants  Products

The arrow is equivalent to an “=“ math. When
we describe the equation we use the word
“yields” or “produces” instead of equals

Example
C + O2  CO2

This reads “carbon plus oxygen react to yield carbon
dioxide”
Balancing a Chemical Equation

A chemical equation is balanced when
the ions or atoms found on the reactant
side of the equation equals that found
on the product side.

The arrow can be considered the
balance point.
Symbols Used in Equations
Solid (s)
 Liquid (l)
 Gas (g)
 Aqueous solution (aq) (dissolved in water)
H2SO4
or
Pt
 Catalyst
 Escaping gas ()
 Change of temperature/ heat energy
( or + 3kJ or – 3kJ) *there is no subtraction…a

negative sign means released/exothermic
Chemical Reactions and
Chemical Equations
Reactants
Products
Driving forces:
a) Color change
b) Formation of a solid/precipitate
c) Evolution of a gas
d) Evolution or absorption of heat
Balancing Equations
– When balancing a chemical reaction you may add
coefficients in front of the compounds to balance the
reaction, but you may not change the subscripts.
 Changing the subscripts changes the compound.
Subscripts are determined by the valence electrons
(charges for ionic or sharing for covalent)
– Think back to naming compounds/ determining
formulas. NaCl exists, because Na is + and Cl is -, but
NaCl2 does NOT exist since you would not have a
neutral compound! You can’t just add a number to a
formula to balance an equation.
Subscripts vs. Coefficients

The subscripts
tell you how
many atoms of a
particular
element are in a
compound. The
coefficient tells
you about the
quantity, or
number, of
molecules of the
compound.
Chemical Equations
4 Al(s) + 3 O2(g) ---> 2 Al2O3(s)
This equation means
4 Al atoms + 3 O2 molecules
---produces--->
2 molecules of Al2O3
AND/OR
4 moles of Al + 3 moles of O2 --produces--->
2 moles of Al2O3
Steps to Balancing Equations
There are four basic steps to balancing a chemical equation.
1. Write the correct formula for the reactants and the
products. DO NOT TRY TO BALANCE IT YET! You must
write the correct formulas first.
**And most importantly, once you write them correctly DO NOT
CHANGE THE FORMULAS!
2.
3.
4.
Find the number of atoms for each element on the left
side. Compare those against the number of the atoms of
the same element on the right side.
Determine where to place coefficients in front of formulas
so that the left side has the same number of atoms as the
right side for EACH element in order to balance the
equation.
Check your answer to see if:
– The numbers of atoms on both sides of the equation
are now balanced.
– The coefficients are in the lowest possible whole
number ratios. (reduced)
Some Suggestions to Help You
Some helpful hints for balancing equations:
 Take one element at a time, working left to right except
for H and O. Metals, then nonmetals are a good way,
too. Save H for next to last, and O until last.
 IF everything balances except for O, and there is no
way to balance O with a whole number, double all the
coefficients and try again. (Because O is diatomic as
an element)
 (Shortcut) Polyatomic ions that appear on both sides
of the equation should be balanced as independent
units
Balancing Equations
___ H2(g) + ___ O2(g) ---> ___ H2O(l)
What Happened to the Other
Oxygen Atom?????
This equation is not balanced!
Two hydrogen atoms from a hydrogen
molecule (H2) combine with one of the
oxygen atoms from an oxygen molecule
(O2) to form H2O. Then, the remaining
oxygen atom combines with two more
hydrogen atoms (from another H2 molecule)
to make a second H2O molecule.
Balance this equation!
Na + Cl2
Na- 1
Cl- 2
NaCl
Na-1
Cl-1
**note that the number of sodiums
balance but the chlorine does not. We
will have to use coefficients in order to
balance this equation.
Inserting subscripts
Na + Cl2
Na- 1
Cl- 2
2 NaCl
Na- 1 2
Cl- 1 2
** Now the chlorine balances but the
sodium does not! So we go back and
balance the sodium.
Finally balanced!
2Na + Cl2
Na- 1 2
Cl- 2
2 NaCl
Na-1 2
Cl-1 2
**Since the number of each element on
the reactant side and the product side of
the equation are equal, the equation is
balanced.
Balancing
Equations
___ Al(s) + ___ Br2(l) ---> ___ Al2Br6(s)
Balancing Equations
Sodium phosphate + iron (III) oxide 
sodium oxide + iron (III) phosphate
Na3PO4 +
Na2O +
Fe2O3 ---->
FePO4
Types of Chemical Reactions (rxns.)
Introduction
– Chemical reactions occur when bonds (between the
electrons of atoms) are formed or broken
– Chemical reactions involve
changes in the chemical composition of matter
(the making of new materials with new properties)
 energy changes

– Bond breaking absorbs Energy (endothermic process)
– Bond making releases Energy (exothermic process)
– Symbols represent elements
– Formulas describe compounds
– Chemical equations describe a chemical reaction
Types of Reactions
•
•
Reactions are classified by their products.
There are five main types of chemical reactions we
will talk about:
1.
2.
3.
4.
5.
•
Synthesis reactions
Decomposition reactions
Single replacement reactions
Double replacement reactions
Combustion reactions
You need to be able to identify the type of reaction
and predict the product(s)
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!
(BrINClHOF) For example, Oxygen is O2 as an
element.
In a compound, it can’t be a diatomic element because
it’s not an element anymore, it’s a compound!
Synthesis Reactions

Synthesis (meaning to make) are
– Also called
 Direct combination or combination reactions
 Addition reactions
– typified by their single product.

If you have a reaction in which at least 2
elements or compounds are reacted and
produce a single product, the reaction is a
synthesis reaction.
Synthesis Reactions
•
reactant + reactant  1 product
Basically: A + B  AB
• Example: 2H2 + O2  2H2O
• Example: C + O2  CO2
• Note: Single Product! This is your clue that this is a synthesis or combination
reaction.
Synthesis Reactions
•
Here is another example of a synthesis
reaction
Examples of Synthesis Reactions

2Na + S
Na2S
– This one is an example of two elements in atomic form
(Na and S) combining to form a compound (sodium
sulfide).

2H2 + O2
2H2O
– In this example, A and B are two elements in molecular
form (hydrogen and oxygen molecules), and the
product is water, which is simply the chemical
combination of hydrogen and oxygen.
Examples of Synthesis Reactions

4Fe + 3O2
2Fe2O3
– In this example, substance “A” is an element in atomic
form (Fe), and substance “B” is an element in
molecular form (O2). The result is a direct chemical
combination of the two elements (Fe2O3, iron (III)
oxide, which is “rust”). * we knew to use iron (III) b/c
3 is the most common ion when we look at the
oxidation state periodic table

CuO + H2O
Cu(OH)2
– This is an example where both substances going into
the reaction are molecules. The result is what you get
when you “add” all of the atoms in the reaction
together.
Practice
•
•
•
•
Predict the products. Write and balance the
following synthesis reaction equations.
Sodium metal reacts with chlorine gas
2 Na(s) + Cl2(g)  2 NaCl(s)
Solid Magnesium reacts with fluorine gas
Mg(s) + F2(g)  MgF2(s) Balanced
Aluminum metal reacts with fluorine gas
2 Al(s) + 3 F2(g) 
2 AlF3(s)
Decomposition Reactions
Decomposition reactions are really just the opposite
of a synthesis reaction. Remember, if you can
make a substance, you should be able to break it
back apart into its components.
A good way to remember decomposition reactions to
to remember what happens when something
decomposes. It falls apart!
Decomposition Reactions
•
•
•
•
•
•
Decomposition reactions occur when a compound
breaks up into the elements or in a few to simpler
compounds
1 Reactant  Product + Product
Basically: AB  A + B
Example: 2 H2O  2H2 + O2
Example: 2 HgO  2Hg + O2
Note: Single Reactant! The single reactant is your clue that this is a decomposition
reaction.
Decomposition Reactions
•
Another view of a decomposition reaction:
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 this class
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) 
Practice
Identify the type of reaction for each of the
following synthesis or decomposition reactions,
and write the balanced equation:
N2 (g) + O2(g)  NO2 (g)
BaCO3(s)  BaO(s) + CO2 (g)
Co(s)+ S(s)  CoS (s)
NI3(s)  N2 (g) +
I2 (s)
Single Replacement Reactions
Single replacement reactions occur when one
chemical takes the place of another in a
reaction.
In the typical single replacement reaction, an
element trades places with one of the ions in
a compound.
Single Replacement Reactions
•
•
•
Single Replacement Reactions:
A metal can replace a metal (+) OR
a nonmetal can replace a nonmetal (-).
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!)
When H2O splits into ions, it splits into
H+ and OH- (not H+ and O-2 !!)
Single Replacement Reactions
•
Another view:
The Activity Series

Not all single replacement reactions will
occur.

This depends upon the location of the
elements present in the activity series

Elements above MAY replace elements
below; elements below MAY NOT replace
elements above them on the series
You will be given a copy of this!!!!
Single Replacement Reactions
•
Write and balance the following single
replacement reaction equation:
• Zinc metal reacts with aqueous hydrochloric
acid
Zn(s) + 2 HCl(aq)  ZnCl2 + H2(g)
Note: Zinc replaces the hydrogen ion in the
reaction
[If ZnCl2 + H2(g)  Zn(s) + HCl(aq) the reaction WOULD NOT
OCCUR because Hydrogen is below zinc on the activity
series]
Single Replacement Reactions
•
Sodium chloride solid reacts with fluorine gas
2 NaCl(s) + F2(g)  2 NaF(s) + Cl2(g)
Note that fluorine replaces chlorine in the compound
•
Aluminum metal reacts with aqueous copper (II)
nitrate
2Al(s)+ 3 Cu(NO3)2(aq) 3Cu(s) + 2Al(NO3)3(aq)
Double Replacement Reactions
•
Double Replacement Reactions occur when a
metal replaces a metal in a compound and a
nonmetal replaces a nonmetal in a compound
• two ions trade places and forming new compounds.
•
•
•
Compound + compound  product + product
AB + CD  AD + CB
Notice that one ion from
compound AB replaces one
ion from compound CD.
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)
Solubility

For a double replacement reaction to have
occurred, a solid (precipitate) MUST be
formed
 There are rules to determine which of the
materials formed is the solid
 If no solid is formed, there is said to be no
reaction.
Figure 8.4: The forming of solid AgCl.
Solubility Tables * see CRH

Solubility tables help determine which materials are soluble in
water and which are not
 In general, Solubility Rules can be summarized as follows
1. All compounds containing alkali metal cations and the ammonium ion
2.
3.
4.
5.
6.
are soluble.
All compounds containing NO3-, ClO4-, ClO3-, and C2H3O2- anions are
soluble.
All chlorides, bromides, and iodides are soluble except those containing
Ag+, Pb2+, or Hg22+.
All sulfates are soluble except those containing Hg22+, Pb2+, Sr2+, Ca2+, or
Ba2+.
All hydroxides are insoluble except compounds of the alkali metals,
Ca2+, Sr2+, and Ba2+.
All compounds containing PO43-, S2-, CO32-, and SO32- ions are insoluble
except those that also contain alkali metals or NH4+.
Practice
•
Predict the products. Balance the equation
1.
HCl(aq) + AgNO3(aq) 
2.
Pb(NO3)2(aq) + BaCl2(aq) 
3.
FeCl3(aq) + NaOH(aq) 
4.
H2SO4(aq) + NaOH(aq) 
Combustion Reactions
Combustion reactions are the ones that burn (or
explode!). There are two types of combustion
reactions—complete or incomplete reactions.
These reactions are identified by their products.
They either produce carbon monoxide and water
or carbon dioxide and water.
Complete Combustion Reactions
These reactions burn “efficiently” which
means they produce carbon dioxide and
water. These reactions typically burn
cleanly and leave very little residue behind.
Combustion Reactions
•
•
Combustion reactions occur
when a hydrocarbon reacts with
oxygen gas.
This 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)
Combustion Reactions
•
In general:
CxHy + O2  CO2 + H2O
• Products in combustion are ALWAYS
carbon dioxide and water. (although
incomplete burning does cause some byproducts like carbon monoxide)
• Combustion is used to heat homes and
run automobiles (octane, as in gasoline,
is C8H18)
Complete Combustion Reactions

CH4 + O2  CO2 + H2O

They may also be written:
O2
CH4  CO2 + H2O
With O2 written above the arrow

Clues: CO2 (carbon dioxide) in the product
along with water
Combustion
•
Example
•
•
C5H12 + 8 O2  5 CO2 + 6 H2O
Write the products and balance the
following combustion reaction:
•
C10H22 + 15.5 O2 10 CO2 + 11 H2O
•
•2 C10H22 + 31 O2 
20 CO2 + 22 H2O
Incomplete Combustion Reactions
Incomplete combustion reactions occur when
something does not burn efficiently. This
can cause a lot of harm if the gases
produced cannot escape. Carbon
monoxide,an odorless and colorless gas, is
dangerous. People poisoned by this gas
usually become sleepy and can die due to
exposure.
Incomplete Combustion Reactions

CH4 + O2  CO + H2O + CO2

These reactions may also be written by:
O2
CH4  CO + H2O + CO2
Again, the O2 is usually written
over the arrow.

Clue: CO (Carbon monoxide as a product.)
Combustion Reactions
Edgar Allen Poe’s
drooping eyes and
mouth are potential
signs of CO
poisoning.
Mixed Practice
•
1.
2.
3.
4.
5.
State the type, predict the products, and
balance the following reactions:
BaCl2 + H2SO4 
C6H12 + O2 
Zn + CuSO4 
Cs + Br2 
FeCO3 
Predicting Products of Reactions

Completing reactions requires knowledge
of the different reaction types (sometimes
called mechanisms).

You must first identify the reaction type
by the reactants. The only type of reaction
that cannot be predicted this way is the
combustion reaction since the products are
very similar.
First Step:

Identify reaction type
Example:
Al + O2 
Clue: 2 elements – Synthesis or combination
reaction
Second Step:

Write the net ionic equation for the
reactants
Al + O2  becomes
Al3+ + O2- 
Step 3

Using clues, complete reaction taking care
to write each formula correctly by checking
charges and “criss-crossing” if necessary.
Al + O2  Al3+O2-
Al + O2  Al2O3
Predicting Products of Reactions (cont.)

For Single Replacement reactions, check
activity series to make sure the reaction
goes.
 Once you write the molecular equation, you
should check for reactants and products that
are soluble or insoluble. (Double
Replacement only)
Reactions in Aqueous Solutions

a.k.a. Net Ionic Equations
 Molecular Equations: shows complete formulas for
reactants and products
– Does not show what happens on the molecular level

Total (or Complete) Ionic Equations: All substances
that are strong electrolytes (are soluble and dissociate) are
written as their ions.
– Some ions participate in the reaction
– Some ions do NOT participate in the reaction-called spectator
ions.

Net Ionic Equations: show only the ions that participate
in the reaction
Writing Total Ionic Equations





Once you write the molecular equation (synthesis,
decomposition, etc.), you should check for
reactants and products that are soluble or
insoluble.
We usually assume the reaction is in water
We can use a solubility table to tell us what
compounds dissolve in water.
If the compound is soluble (does dissolve in
water), then splits the compound into its
component ions
If the compound is insoluble (does NOT dissolve
in water), then it remains as a compound
Writing Total Ionic Equations
Molecular Equation:
K2CrO4 + Pb(NO3)2  PbCrO4 + 2 KNO3
Soluble
Soluble
Soluble
Insoluble
Total Ionic Equation:
2 K+ + CrO4 -2 + Pb+2 + 2 NO3- 
PbCrO4 (s) + 2 K+ + 2 NO3-
Net Ionic Equations

These are the same as total ionic equations, but
you should cancel out ions that appear on
BOTH sides of the equation
Total Ionic Equation:
2 K+ + CrO4 -2 + Pb+2 + 2 NO3- 
PbCrO4 (s) + 2 K+ + 2 NO3(Spectator ions)
Net Ionic Equation:
CrO4 -2 + Pb+2  PbCrO4 (s)
Net Ionic Equations

Try this one! Write the molecular, total ionic, and net ionic
equations for this reaction: Silver nitrate reacts with Lead
(II) Chloride in hot water.
AgNO3 + PbCl2 
Molecular:
2 AgNO3 + PbCl2  2 AgCl + Pb(NO3)2
Total Ionic:
2 Ag+ + 2 NO3- + Pb+2 + 2 Cl-  2 AgCl (s) + Pb+2 + 2 NO3Net Ionic:
2Ag+ + 2Cl-  2AgCl (s)
Acid-Base Reactions

Acid:
– produces hydrogen ions (H+) in solution (Arrhenius)
– proton donor (Lewis)

Base:
– produces hydroxide ions (OH-) in solution (Arrhenius)
– proton acceptor (Lewis)

The reaction ALWAYS forms water and an ionic
compound (mostly aqueous, known as a salt).
–
The actual definition of a salt is the ionic product of an acid-base
neutralization reaction)
Acid-Base Neutralization

Chemically the reaction looks like this:
Acid + Base
Salt + Water
 A classic example:
HCl + NaOH
NaCl + H2O
Hydrochloric Acid
Water
Sodium Hydroxide (Lye)
Sodium Chloride
(Table Salt)
Acid-Base Reactions

Example
HNO3 (aq) + KOH (aq) 
Molecular:
H2O (l) + KNO3 (aq)
HNO3 (aq) + KOH (aq) 
Total Ionic:
H+ (aq) + NO3- (aq) + K+ (aq) + OH- (aq)
 H2O (l) + K+ (aq) + NO3- (aq)
Net Ionic:
H+ (aq) + OH- (aq)  H2O (l)
Acid-Base Neutralization

Here’s the equation again:
– HCl + NaOH
NaCl + H2O
– Chemically, this is a double replacement
reaction:


The H traded its Cl for an OH
The Na traded its OH for a Cl.
Oxidation-Reduction Reactions

a.k.a. Redox Equations
 Between a metal and a nonmetal forming an
ionic compound
– Electron transfer occurs
– Oxidation numbers: assigning an excess or
deficiency in electrons for each element (the
charge based on the compound).
Rules for Oxidation Number (ox. #)
Determination
1. The sum of the oxidation numbers add up to the
charge
a. all elements have an ox. # of 0
b. ions of elements, ox. # is the charge (Cl-)
c. the sum of the ox. # of a complex ion equals the
charge (CO3-2 )
2. H is 1+ when combined with a nonmetal and 1with a metal
H3PO4
H= 1+ PO4 = 3-
CaH2
Ca = 2+ H= 1-
Rules for Oxidation Numbers (cont.)
1. F is always 1-; Cl, Br, I are 1- except when
combined with each other or O
2. O is 2- except when combined with F (F2O)
3. Group I is 1+ and Group II is 2+ in their
compounds
Recognizing Redox Rxns.
oxidation

2+ 2(-1)
1+ 1-
2 HCl (aq) + Mg 0(s)  MgCl2 (aq) + H20 (g)
reduction

Net:
2 H+ (aq) + Mg0 (s)  Mg 2+ (aq) + H2 0 (g)


Loss of electron = oxidation
Gain of electron = reduction
– “LEO the lion goes GER"
Half Reactions

Separate the individual oxidation and reduction
reactions.
 Look at electron movement
 Half rxn.:
Mg0 (aq)  Mg 2+ + 2 e2e- + 2 H+  H20

Net:
2 H+ (aq) + Mg0 (aq)  Mg 2+ (aq) + H2 0 (q)

Oxidizing agent: the one reduced (H+)
 Reducing agent: the one oxidized (Mg0)
Recognition of Redox rxns.



Oxidation # changes
Reactions with oxygen
Reaction of any element (forms a new
compound)
Balancing



Balance by mass
Balance by charge
Balance net ionic equation
Example Problem:
Fe (s) + Cl2 (aq)  FeCl3 (aq)
1.
Balance by mass
2 Fe (s) + 3Cl2 (aq)  2 FeCl3 (aq)
2.
Write the ionic equation
2 Fe (s) + 3 Cl2 (aq)  2 Fe3+ (aq) + 6 Cl- (aq)
3.
Write half reaction
2 Fe0 (s)  2 Fe 3+ + 6 e6e- + Cl2  2 Cl-
3. Balance by charge (want # of e- to cancel)
2 Fe0 (s)  2 Fe 3+ + 6e6e- + 3Cl20  6 Cl+
2 Fe0 (s) + 6e- + 3Cl2  2 Fe 3+ + 6e- + 6 ClFinal eqn.:
2 Fe (s) + 3Cl2  2 Fe 3+ + 6 Cl-