Stoichiometry
CHAPTER 12
The Arithmetic of Equations
SECTION 1
SECTION 1 LEARNING TARGETS
12.1.1 – I can explain how balanced equations
apply to both chemistry and everyday life.
12.1.2 – I can interpret balanced chemical
equations in terms of moles, representative
particles, mass, and gas volume at STP.
12.1.3 – I can identify the quantities that are
always conserved in chemical reactions.
USING EVERYDAY EQUATIONS

A balanced chemical equation provides the
same kind of quantitative information that a
recipe does.
USING BALANCED CHEMICAL EQUATIONS

Chemists use balanced chemical equations as
a basis to calculate how much reactant is
needed or product is formed in a reaction.
EXAMPLE:

Tiny tike has decided to make 288 tricycles
each day. How many tricycle seats, wheels,
and pedals are needed?

Stoichiometry – the calculation of quantities in
chemical reactions.
INTERPRETING CHEMICAL EQUATIONS

A balanced chemical equation can be
interpreted in terms of different quantities,
including; numbers of atoms, molecules, or
moles, mass, and volume.
NUMBER OF ATOMS

A balanced equation indicates that the number
and type of each atom that makes up each
reactant also makes up each product.
NUMBER OF MOLECULES
Here the coefficients tell you how many
molecules will react and form.
 Much like how many atoms.

MOLES
A balanced equation also tells you the number
of moles of reactants and products.
 The coefficients tell you this.
 You’ll use this most often.

MASS
A balanced equation obeys the law of
conservation of mass.
 Using a mole relationship you can relate
number of moles to mass of either reactants or
products.

VOLUME
If you are at STP the equation tells you about
volumes of gases.
 You can use mole relationships for this also.

MASS CONSERVATION IN CHEMICAL REACTIONS

Mass and atoms are conserved in every
chemical reaction.
EXAMPLE:
Balance the following equation.
___C2H4(g) + ___O2(g) → ___CO2(g) + ___H2O(g)
 Interpret the balanced equation in terms of
relative numbers of moles, volumes of gas at
STP, and masses of reactants and products.

Chemical Calculations
SECTION 2
SECTION 2 LEARNING TARGETS
12.2.1 – I can construct mole ratios from
balanced chemical equations and apply these
ratios in stoichiometric calculations.
12.2.2 – I can calculate stoichiometric quantities
from balanced chemical equations using units
of moles, mass, representative particles, and
volumes of gases at STP.
WRITING AND USING MOLE RATIOS

Mole ratio – a conversion factor derived from
the coefficients of a balanced chemical
reaction interpreted in terms of moles.

In chemical calculations, mole ratios are used
to convert between moles of reactant and
moles of product, between moles of products,
or between moles reactants.
EXAMPLE:
4Al(s) + 3O2(g) → 2Al2O3(s)
 Write six mole ratios that can be derived from
this equation.

MOLE-MOLE CALCULATIONS
The easiest way to see these is to do an
example.
 W is the unknown, G is the given quantity.
 a and b are the coefficients from the balanced
chemical equation.

4Al(s) + 3O2(g) → 2Al2O3(s)
 How many moles of aluminum are needed to
form 3.7 moles of aluminum oxide?

MASS-MASS CALCULATIONS
In a reaction things are not measured in moles
(no scale does this).
 Instead things are measured in mass then
transferred to moles.

STEPS IN SOLVING A MASS-MASS PROBLEM
1.
Change the mass of G to moles of G (mass G
→ mol G) by using the molar mass of G.
2.
Change the moles of G to moles of W (mol G
→mol W) by using the mole ratio from the
balanced equation.
3.
Change the moles of W to grams of W (mol W
→ mass W) by using the molar mass of W.
EXAMPLE:
Acetylene gas (C2H2) is produced by adding
water to calcium carbide (CaC2).
CaC2(s) + 2H2O(l) → C2H2(g) + Ca(OH)2(aq)
 How many grams of acetylene are produced by
adding water to 5.00g of calcium carbide?

OTHER STOICHIOMETRIC CALCULATIONS
In a typical stoichiometric problem, the given
quantity is first converted to moles.
 Then the mole ratio from the balanced
equation is used to calculate the number of
moles of the wanted substance.


Finally, the moles are converted to any other
unit of measure related to the unit mole, as the
problem requires.
EXAMPLE:

How many molecules of oxygen are produced
by the decomposition of 6.54g of potassium
chlorate (KClO3)?
2KClO3(s) → 2KCl(s) + 3O2(g)
EXAMPLE:
The equation for the combustion of carbon
monoxide is: 2CO(g) + O2 → 2CO2(g)
 How many liters of oxygen are required to burn
3.86L of carbon monoxide?

Limiting Reagent and Percent Yield
SECTION 3
SECTION 3 LEARNING TARGETS
12.3.1 – I can identify the limiting reagent in a
reaction.
13.3.2 – I can calculate theoretical yield, actual
yield, or percent yield given appropriate
information.
LIMITING AND EXCESS REAGENTS
When cooking you know you need the right
amounts of ingredients for the recipe to turn
out.
 In a chemical reaction, an insufficient quantity
of any of the reactants will limit the amount of
product that forms.


Limiting reagent – determines the amount of
product that can be formed in the reaction.
 This
is the one used up first in a reaction.
 The reaction can only “go” until this reactant is
completely used up.

Excess reagent – reactant that is not
completely used up in a reaction.
How would the amount of products formed if you started with
four molecules of N2 and three molecules H2?
EXAMPLE:
The equation for the complete combustion of
ethene (C2H4) is:
C2H4(g) + 3O2(g) → 2CO2(g) + 2H2O(l)
 If 2.70mol of C2H4 is reacted with 6.30mol of
O2, identify the limiting reagent.


The reactant that is present in the smaller
amount by mass or volume is not necessarily
the limiting reagent.
EXAMPLE:
The heat from an acetylene torch is produced
by burning acetylene (C2H2) in oxygen:
2C2H2 + 5O2 → 4CO2 + 2H2O
 How many grams of water can be produced by
the reaction of 2.4 mol C2H2 with 7.4 mol O2?

PERCENT YIELD

Your grades are usually expressed as a percent
 right

 100 % 

 total

Chemists use similar calculations when
products are formed based on balanced
equations.
 In theory all reactions would produce at 100%.
 In reality they don’t.

Theoretical yield – the maximum amount of
product that could be formed from given
amounts of reactants.
 Actual yield – the amount of product that
actually forms when the reaction is carried out.


Percent yield – ratio of the actual yield to the
theoretical yield expressed as a percent.
The percent yield is a measure of the efficiency
of a reaction carried out in the laboratory.
 Percent yield can be lowered by:

 Impure
reactants.
 Loss of product in filtration or transferring.
 If reactants or products have not been carefully
measured.
EXAMPLE:
When 84.8g of iron (III) oxide reacts with an
excess of carbon monoxide, iron is produced.
Fe2O3(s) + 3CO → 2Fe(s) + 3CO2(g)
 What is the theoretical yield of iron?

EXAMPLE:
If 50.0g of silicon dioxide is heated with an
excess of carbon, 27.9g of silicon carbide is
produced. SiO2(s) + 3C(s) → SiC(s) + 2CO(g)
 What is the percent yield of this reaction?
