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Balancing Chemical Equations
Balanced Equations
• Atoms can not be created or destroyed
• All atoms we start with we must end up
with…and vice versa!
• A balanced equation has the same
number of each element on both sides of
the equation.
• C + O2  CO2
• This equation is already balanced
• But what if an equation isn’t already
balanced?
Like this….
• C + O2  CO
• We need one more ________ in the
products.
• You can’t change the formula, because it
describes what is being produced…CO
(Carbon Monoxide.)
This gives a better/closer idea of what is happening…
BUT….
• The other oxygen must be used to make
another CO
• But where does the other C come from?
• Must have started with two C’s
• 2 C + O2  2 CO
Rules for Balancing
• Write the correct formulas for all the
reactants and products
• Count the number of atoms of each type
appearing on both sides.
• Balance the elements one at a time by
adding coefficients (the numbers in front)
2 CO2
• Check to see if it is balanced
Never
• Never change a subscript to balance an
equation CO2
– If you change the formula you are describing
a different reaction.
– H2O is a different compound than H2O2
• Never put a coefficient in the middle of a
formula
– 2 NaCl is ok Na2Cl is not.
Example
H2 + O2  H2O
• Make a table to keep track of atoms
Example
H2 + O2  H2O
R
P
2
H
2
2
O
1
Need twice as much O in the product
Example
H2 + O2  2H2O
R
Changes the O
P
2
H
2
2
O
1
Example
H2 + O2  2H2O
R
Also changes the H
P
2
H
2
2
O
1
2
Example
H2 + O2  2H2O
R
P
2
H
2
4
2
O
1
2
Now we need twice as
much H in the reactant
Example
2H2 + O2  2H2O
R
Recount to check
P
2
H
2
4
2
O
1
2
Example
2H2 + O2  2H2O
Your answer
R
4
Recount to check
P
2
H
2
4
2
O
1
2
Types of Reactions
•
•
•
•
Millions of reactions
Too many to remember
They fall into several categories
We will focus on Double Replacement in
today’s lab
Double Replacement
• Two things replace each other
• Reactants must be two ionic compounds or
acids.
• Usually in aqueous solution
• NaOH + FeCl3 
• The positive ions change place
• NaOH + FeCl3  Fe+3OH- +Na+1Cl-1
• NaOH + FeCl3  Fe(OH)3 + NaCl
Double Replacement
• Will only happen if one of the products
– Doesn’t dissolve in water and forms a solid
• (look at solubility rules)
– Or is a gas that bubbles out
– Or is a covalent compound usually water
Potassium iodide
After adding lead nitrate
2KI(aq) + Pb(NO3)2 (aq)  2KNO3(aq) + PbI2 (s)
PbI2 lead (II) iodide
is insoluble
General Rules for the Water Solubilities
of Common Ionic Compounds
• Compounds that are mostly soluble:
– All nitrates
– Alkali metal (group 1A) and ammonium
compounds
– Chlorides, bromides, and iodides, except for
those of Pb2+, Ag+, Hg2+
– Sulfates except for those of Sr2+, Ba2+, Pb2+,
and Hg2+
• CaSO4 is slightly soluble
General Solubility Rules
• Compounds that are mostly insoluble:
– Carbonates, hydroxides, and sulfides, except
for ammonium compounds and those of the
group 1A metals. (The hydroxides and
sulfides of Ca2+, Sr2+, and Ba2+ are slightly to
moderately soluble.)