Chemistry
Chapter 12
Notes #3
Limiting and Excess
Reactions
• What causes a chemical reaction to stop?
• The reactants – or one of the reactantsget used up.
• Limiting reactant – the reactant that gets
used up first, therefore limiting (stopping)
the reaction
Example
• How many sets of blue/red/yellow crayons
can you make with the following?
– 10 red crayons
– 12 blue crayons
– 6 yellow crayons
• What is the limiting reactant?
• Which were in excess (those not used
up)?
Example
Hydrogen
+
Oxygen
4 Water molecules get created and an oxygen molecule is left over (excess),
which means that Hydrogen is the limiting reactant
The Math!
• CsF + XeF6 -> CsXeF7
• How many moles of cesium xenon
heptafluoride can be produced from the
reaction of 12.5 mol cesium fluoride with
10.0 mol of xenon hexafluoride?
• Hint – notice the 1:1:1 ratio
• 10 mol CsXeF7
• You can only make as much as your least
amount !
The Math
• 6Na + Fe2O3 -> 3Na2O + 2Fe
• If 100.0 g Na and 100.0 g Fe2O3 are used,
determine the limiting reactant
•
•
•
•
•
1) reactants = grams to moles
2) Mole to Mole comparison (ratio)
3) Limiting reactant becomes “given”
4) Go to Stoichiometry step #3
Excess reactant: Start Mass – Mass used
The Math
• 6Na + Fe2O3 -> 3Na2O + 2Fe
• If 100.0 g Na and 100.0 g Fe2O3 are used,
determine the limiting reactant
The Math
• 6Na + Fe2O3 -> 3Na2O + 2Fe
• What is the mass of solid iron produced?
The Math
• 6Na + Fe2O3 -> 3Na2O + 2Fe
• What is the mass of the excess reactant
after the reaction?
Percent Yield
• If the conditions aren’t right, often a
reaction will stop before the reactants are
used up.
• Percent yield shows the efficiency of the
reaction. (compare to a letter grade on a
test)
• Actual yield (experiment) x 100
• Theoretical yield (math)
Percent Yield
• How might a business use percent yield in
production for overall cost effectiveness?
• Overall, you want percent yield to be as
close to 100% as possible
– Might result in changing conditions
– Or more excess (which speeds up the
reaction or help make it go through to
completion)
The Math
• 6Na + Fe2O3 -> 3Na2O + 2Fe
• If we ran this experiment and produced
62.0 g of Fe, what would the percent yield
be?