Chemistry 521
Chapter 4
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Stoichiometry is the study of the relative
quantities of reactants and products in
chemical reactions.
Ex:
Two hydrogen molecules and one oxygen
molecule produce two water molecules.
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Stoichiometry is also the ratio of atoms
and/or molecules as shown by the chemical
equation.
Recall the Law of Definite Proportions – atoms
combine in definite fixed proportions.
So then the chemical equation is also a ratio:
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Consider the analogy on Page 111
◦ 3 slices of toast + 2 slices of turkey + 4 strips of
bacon = 1 sandwich
◦ 3:2:4:1 ratio
The ratio can be multiplied.
◦ 6:4:8:2
◦ 9:6:12:3
The equation provides very useful information.
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Consider the equation to produce ammonia.
1 molecule N2: 3 molecules H2 : 2 molecules
NH3
You can multiply this ratio:
◦ By 2:
◦ 2 molecule N2: 6 molecules H2 : 4 molecules NH3
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By 10:
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10 molecule N2: 30 molecules H2 : 20 molecules NH3
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By 2.63 × 1014:
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2.63 × 1014 molecule N2: 7.89 × 1014 molecules H2 : 5.26 ×
1014 molecules NH3
Every chemical equation ratio always holds
true no matter what you multiply by.
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Suppose you wanted 20 molecules of
ammonia, how many molecules of N2 are
required? H2?
The chemical equation is:
Therefore the ratio is:
1N2 : 3H2 : 2NH3
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To get 20 NH2, multiply everything by 10
10 N2 : 30 H2 : 20 NH3
OR create conversion factors:
20 molecules NH3 × 1 molecule N2 = 10 molecules N2
2 molecules NH3
20 molecules NH3 × 3 molecules H2 = 30 molecules H2
2 molecules NH3
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Do Practice Problem #1 together on Page 114
Do Practice Problems on Page 114 #s 2 & 3
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Consider:
As we have said the coefficients represent a ratio
of molecules
1 molecule N2: 3 molecules H2 : 2 molecules NH3
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These coefficients from the chemical equation
would also represent the number of moles of
each atom or molecule.
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6.02 × 1023 (1 molecule N2: 3 molecules H2 :
2 molecules NH3)
Would still give the same ratio molecule to
molecule.
Mole Ratios are the relationships between
moles in a balanced equation.
Note- each equation about mole ratios may
only involve 2 components of the chemical
equation. It is assumed that the other
atoms/molecules will follow in the same
trend and that there is enough of each
reactant.
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Do Practice Problem #4 as an example on
Page 11
Do Practice Problems on Page 115 #s 5-7
Mole to Mole Worksheet
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Recall the Law of Multiple Proportions – that
when 2 elements combine two or more
different compounds may result – depending
on the conditions and the amount of
reactants available.
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Ex: carbon and oxygen can combine to form
carbon dioxide or carbon monoxide.
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Do Sample Problem on Page 116
Do Practice Problems on Page 117 #s 8-10
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In Chapter 2 you learned that once you have a
number in moles you could convert to grams,
volume, number of particles or vice versa.
The coefficients in a chemical equation are
considered the number of moles.
Therefore, if you have two moles of N2, you
can calculate the mass of N2 in that reaction.
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Consider:
◦ 1 mol of N2 × 28.02 g = 28.02 g N2
1 mol
◦ 3 mol of H2 × 2.02 g = 6.06 g H2
1 mol
◦ 2 mol of NH3 × 17.04 g = 34.08 g NH3
1 mol
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According to the Law of Conservation of
Matter – matter is neither created nor
destroyed, it only changes form
Therefore the mass of the reactants is equal
to the mass of the products in a chemical
equation which is the Law of Conservation of
Mass
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Whether you add up the mass of the reactants
or the products, you will always get the same
result.
N2 + 3H2
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2NH3
28.02g + 6.06 g
2(14.01 +3(1.01)) g
= 34.08 g
= 34.08 g
Note – for stoichiometric problems the
equations MUST be balanced.
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If you know the quantity of one substance in
a chemical reaction (in particles, moles,
grams, or liters), you can calculate the
quantity of any other substance in the
reaction.
Stoichiometry is so useful because you can
use it to make predictions.
One purpose of stoichiometric calculations is
to determine how much of a reactant is
needed to carry out a reaction.
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Stoichiometric analysis involving mass is
called gravimetric stiochiometry.
Stoichiometric calculations involving volume
of gases is called gas stoichiometry.
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Do Sample Problem on Page 120
Do Practice Problems on Page 120 #s 11-14
Do Sample Problem on Page 121
Do Practice Problems on Pages 122-123 #s
15-18
1.
2.
3.
4.
Write the balanced equation.
Convert the number given in the problem to
moles (use Mole Chart from Ch. 2).
Convert between atoms/molecules using
mol ratio based on the chemical equation.
Convert this mole value to what was asked
for in the question.
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Do Sample Problem on Page 124
Do Practice Problems on Page 125 #s 19-22
Do Section Review on Pages 126-127
Worksheet
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Recall the clubhouse sandwiches from the
previous section. By knowing the base formula
you could predict how much of each ingredient
you need to make multiple sandwiches.
3 slices of toast + 2 slices of turkey + 4 strips of bacon = 1 sandwich
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What if you had 6 slices of toast, 12 slices of
turkey, and 20 strips of bacon? How many
sandwiches could you make?
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Taking all the ingredients you have into
account and the formula, since you only have
6 slices of toast, you could only make 2
complete sandwiches.
There would be 8 slices of turkey and 12
strips of bacon left over or extra.
The toast would be the limiting reactant in
the sandwich equation.
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The limiting reactant is the one to run out
first.
Once the toast was used up you were unable
to make any more sandwiches.
The slices of turkey and strips of bacon would
be left over or in excess.
The excess reactant is the one that is left
unreacted.
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Do Thought Lab on Page 129
The concept of limiting reactant and excess
reactant come up in real life situations.
Zinc is extracted from zinc oxide by adding
carbon.
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In industrial settings, they do not try to
ensure the reactants are in a 1:1 ratio. They
just make sure they have more carbon
(charcoal) than zinc oxide.
This reaction will produce zinc until one
reactant runs out and to ensure the most
profit they do not want to waste any zinc
oxide.
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Often you have to identify which of the
reactants is the limiting reactant.
One method is to predict how much product
each reactant could, at most, produce.
The reactant that produces the least amount
of product is the limiting reactant.
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Do Sample Problem on Page 130
Do Practice Problems on Page 131 #s 23 - 26
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Stoichiometry is useful because it can make
predictions.
Chemists need to know how much product
they can expect from a certain reaction.
Often experimental results are compared to
the calculated predicted results.
By analyzing an impure substance with a
known substance the predicted expected
mass is compared to the actual mass of the
substance.
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Since most chemical reactions are not
balanced inside a beaker, you must first
determine the limiting reactant to solve
stoichiometric problems.
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Do Sample Problem on Page 133
Do Practice Problems on Pages 134-135 #s
27-30
Do Section Review on Pages 135-136
Limiting Reactant Worksheet
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When you get a mark on a test, how was that
mark calculated? Also recall mass percent.
Your mark × 100%
Total Marks
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The stoichiometric calculations show us the
theoretical yield. This would be how much
product is produced under ideal conditions.
In the real lab, the experiment may not
produce as much product as you would have
expected. This is the actual yield.
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Why does the actual yield differ from the
theoretical yield?
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Competing reactions
Experimental design and technique
Impure reactants
Faulty measuring devices
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These 4 factors lead to inaccurate results.
The accuracy of an experiment is measured by
how close your result is to the expected value.
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The precision of the results refers to the
reproducibility of the results. The results
were precise if they can be repeated again
and again.
Precision depends on the measuring devices,
as in how many decimal places they go to and
your ability to use the devices properly.
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Competing reactions can occur during
experiments.
Ex: phosphorus and chlorine gas forming
phosphorus trichloride. Sometimes the PCl3
then reacts with some of the Cl2 and another
product, PCl5 is formed.
Therefore the yield of the product, PCl3, is
less than calculated.
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The percentage yield is telling you how close
you are to the expected yield.
First find the limiting reactant and use it to
find the amount of product. This is the
theoretical yield. The actual yield will be
given in the problem.
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Do Sample Problem on Page 138
Do Sample Problem on Page 138 #31
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Do Practice Problems on Page 139 #s 32 & 33
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Using the same equation, if the percentage
yield is known and the theoretical yield can
be calculated, then chemists can predict the
actual yield from a sample.
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Do Sample Problem on Page 140
Do Practice Problems on Page 141 #s 34-37
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The percentage yield of a reaction is
necessary in the pharmaceutical industry.
Pharmaceutical companies developing new
drugs may only get a small percentage yield
of product.
In the R&D phase that is reasonable, as
scientists are working on developing a
process or perfecting a reaction.
But the higher the percentage yield then more
money is made.
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The lower the percentage yield the more
costs are associated.
Once a new drug is ready, it would then be
manufactured in large quantities (more cost
effective).
At this point a small difference in percentage
yield can cost companies thousands of
dollars.
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When doing stoichiometric calculations, we
assume ideal conditions.
However, in reality the reactants are not pure
and this can lower the actual yield.
Ex. A 1.00 g sample of sodium chloride
(NaCl) may have absorbed some water. So
the sample is not 1.00 g of just NaCl. You
really do not know how much NaCl you
actually have.
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The percentage purity of a sample tells you
how much of the sample is the element of
compound that you are looking for.
◦ This concept is used in mining situations.
◦ Ore is mined but is not pure.
◦ There is often rock, dirt, or other waste within a
sample.
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To find percentage purity, calculate the
amount of “important” element or compound
in the sample.
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This is the theoretical mass and use it to find
the percentage purity.
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Do Sample Problem on Page 146
Do Practice Problems on Page 147 – 148 #s
38 – 40
Do Section Review on Page 148 #s 1-3
Chapter 4 Review on Pages 149 – 150 #s 1 –
10, 14 - 21