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Stoichiometry II: Molar Mass
PRE-LAB ASSIGNMENTS:
Look up (elements) or calculate (compounds) the mass of one mole of each of the substances
listed in the data table. Show your work and report your results clearly on a piece of 8.5”×11”
notebook paper to be turned in at the start of lab. Enter these values in the Molar Mass column
of Data Table 1 in the Lab Report.
STUDENT LEARNING OUTCOMES:



Describe the amount of substance in terms of the number of units (atoms or molecules),
in addition to mass and/or volume
Convert between number of grams of substance and moles, and vice versa.
Determine the amount of an element, in units of mass or moles, that is present in a
sample of a compound
EXPERIMENTAL GOALS:
The goal of this experiment is to measure out approximately 1 mole amounts of several
substances, make observations, and compare the masses, volumes, and numbers of formula units
in these 1 mole samples.
INTRODUCTION:
In the chemistry laboratory, the amount of a substance is most frequently measured in
terms of mass or volume. The masses and volumes of atoms of different elements are not the
same, and compounds are comprised of atoms of different elements combined in widely varying
ratios. As a result, mass and volume measurements do not give us a direct measure of the
composition of a sample or substance. The mole is a unit for the amount of a substance that
does give us direct insight into atomic/ionic/molecular composition. Chemists use the mole
(sometimes abbreviated as “mol”)1 to designate the amount of a substance in the same way that a
baker may designate 12 eggs as one dozen eggs. Just like the unit of one dozen, the mole is a
standard quantity that can be used as a unit itself (a mole of carbon atoms). Furthermore, the
mole specifies a number, called Avogadro’s number (NA = 6.0221023), in much the same way
that a dozen specifies exactly 12 units. Hence we can say that one mole of carbon atoms is
equivalent to 6.0221023 carbon atoms. The concept of the mole will be useful in much of the
rest of chemistry that you encounter in this course and beyond.
The number below the symbol of an element on a typical periodic table is the value of
both the atomic weight and the molar mass of the element. The atomic weight is the average
1
“Mole” should never be abbreviated as “m” or “M,” since those abbreviations have other meanings.
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mass of an atom of the element in a terrestrial (found on the earth) sample of the
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element measured in atomic mass units (u, 1 u ≈ 1.6605×1024 g). The molar mass
is the mass of one mole of a terrestrial sample of that element. For example, the
atomic weight of chlorine in the example to the right is 35.453 u and the molar mass
35.453
is 35.453 g/mol. The two quantities, atomic weight and molar mass are numerically
the same because both are defined in terms of the mass of carbon-12 atoms. This
means that one mole of chlorine atoms (≈6.022141023 chlorine atoms) has a mass of 35.453 g.
Similarly, one mole of hydrogen atoms would have a mass of 1.00797 g, and a mole of iron
atoms would have a mass of 55.847 g.
Cl
The molar mass of a compound (sometimes referred to as the “molecular weight” for
molecular compounds and the “formula weight” for ionic compounds) is obtained by adding up
the molar masses of all of the elements in the compound multiplied by the number of atoms of
each element in the compound. For example, the molar mass of ammonia, NH3, is obtained by
adding the molar mass of nitrogen and three times the molar mass of hydrogen:
Molar mass of NH3 = 1  molar mass of N + 3  molar mass of H
= 1  14.0067 g/mol + 3  1.00797 g/mol
= 17.031 g/mol
Definition of Formula Unit. In this lab, we will want to think about how many “units” of
our compound are in a mole of the compound. Consider the phrase: “one mole of water.” What
does this mean? Since water is made up of water molecules, it seems reasonable that we have
one mole (or Avogadro’s number) of water molecules. However, when we talk about an ionic
compound, we cannot accurately say that 1 mole of magnesium chloride contains one mole of
magnesium chloride molecules, because magnesium chloride is not made up of molecules, but of
magnesium ions and chloride ions in a 1:2 ratio. To clarify this we can refer to a formula unit of
a compound. A formula unit is the smallest collection of atoms that can have all of the
properties of a compound and is represented by that compound’s formula. When we talk about
having a mole of a substance, we mean that we have one mole of formula units of that substance
in the sample.
Just like a dozen bicycles might be made up of 2 dozen wheels and one dozen frames, a
mole of formula units can have different numbers of moles of the atoms or ions that make up the
compound. For example, a mole of water molecules contains two moles of hydrogen atoms and
one mole of oxygen atoms. Similarly, sodium sulfate (Na2SO4) contains two moles of Na+ ions
and one mole of SO42- ions; it also contains 1 mole of sulfur atoms and 4 moles of oxygen atoms.
The mole is one of the most useful concepts in chemistry due to the fact that ratios betweens
chemical entities in a compound or chemical species is exactly matched by the mole ratio of
those species. The mole provides us with a way of measuring macroscopic amounts of
elements, ions or compounds and relating that information to chemical formulas and equations.
The mole provides us with a method of relating macroscopic measurements to what is occurring
at the atomic/molecular level!
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PROCEDURE:
Prelab Assignment: Make sure that you have completed the Pre-Lab Assignment described at
the beginning of this chapter before coming to lab.
A. General Directions and Setup
Students will work in pairs. Each pair will measure out approximately one mole of one of the
substances listed in the data table, to within ± 0.2 g of the Molar Mass. The sample will be
placed in a graduated cylinder of an appropriate size to estimate the volume of a sample to at
least two significant figures. A label should be placed on that cylinder that gives the name and
formula of the substance and the actual mass of the sample. The samples thus prepared will be
placed in one location in the lab, and each student will use this display to complete Data Table A.
Your instructor will assign each pair of students a substance, or will guide you in the choosing a
substance.
B. Obtaining and Comparing ~1 Mole of Various Substances
Most of the balances in our labs have a 120 or 150 gram capacity. A 100 mL beaker weighs ca.
44 g. If the weight of your sample and the beaker taken together is greater than 120 g, you will
have to use a special balance located in one of the labs to weigh your sample: consult your
instructor for further directions. For sucrose, use a 400 mL beaker for weighing. Use a 250 mL
beaker for benzoic acid, methyl salicylate, and magnesium nitrate hexahydrate. Use a 100 mL
beaker for all other substances.
High accuracy and precision are not required for this lab. Therefore we will use the “tare”
function of the balances to speed the weighing of samples. To use the tare function, place the
beaker on the balance and push the tare or zero button. The scale should read 0.000 g (or 0.0 g
on the high capacity scale). Add substance to the beaker until you get a mass that is within 0.2 g
(within 1 g if your sample will be greater than 100 g) of the molar mass. Record the mass of the
sample directly from the scale to as many significant digits as the balance permits. (It is possible
that the scale zero has drifted since the balance was “tared” (zeroed) so that this mass may be
slightly off. This will not affect the results in this lab, but the tare method is not recommended
when accuracy is required.)
Use the volume level in the beaker to estimate a graduated cylinder size to transfer your sample
into. We will use the graduated cylinder to approximate the volume of the sample to at least 2
significant digits. Transfer the sample to the graduated cylinder and record the volume. For
solids, do not try to estimate the volume to 1/10 of the smallest division. Record this volume in
your data sheet.
Use the mass of your sample to calculate the actual number of moles of substance you measured
and the number of formula units in your sample. You and your partner(s) should do the
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calculations independently and then compare your results. On a piece of labeling tape, record
the following information:
Name of sample
Grams of sample
Approx. volume of sample
Moles of sample
Number of formula units of sample
Also record this information on Data Table 1.
Place the label on your cylinder and then place you cylinder on the designated table or counter so
that the sample can be available for other students to compare.
C. One Mole of Gas
Examine the volume of a mole of gas in the sample container which is provided. Note that a
mole of any gas will have this volume under identical pressure and temperature conditions.
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LAB REPORT
Stoichiometry II: Molar Mass
Name ________________________________
Date _________
Partner ________________________________
Section _________
Report Grade ______
D. Data Analysis.
Record your observations on the molar amounts of various substances in Data Table 1. Use this
information to answer the questions below.
1. In general, will one mole of a substance have the same mass as a mole of a different
substance? Why is this so?
2. In general, do one mole quantities of different substances have the same volume? Why do
you think this is so?
3. The apparent volume of your sample was not always the same for one mole of some of the
samples looked at in the study. What caused the variation? Under what conditions would it
be reasonable to use volume to measure the amount of a substance? Under what conditions
should you be cautious using volume for that purpose?
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4. What property did all of the one mole samples have in common?
5. For gases, one mole of almost any gas will have the same volume under the same conditions
of temperature and pressure. Think about the behavior of the species (molecules, atoms or
ions) that make up solid, liquid and gas samples and try to explain why one mole of different
gasses may have the same volume, but molar quantities of solids and liquids rarely have the
same volume.
E. Study Problems. Use the information in Table 1 to answer the following questions.
1. Determine the amount of each of the following entities there are in 1 molecule of SiO2:
_______ total atoms
_______ O atoms
_______ Si atoms
2. Determine the amount of each of the following entities that are in 1 dozen water molecule?
_______ total atoms
_______ O atoms
_______ H atoms
_______ dozens of atoms
_______ dozens of O atoms
_______ dozens of H atoms
3. Determine the amount of each of the following entities that are in 1.5 dozen sugar molecules?
_______ dozens of total atoms
_______ dozens of H atoms
_______ C atoms
_______ O atoms
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F. Questions.
For the questions below, report your answers to the correct number of significant figures.
1. How many moles of oxygen atoms were in your
group’s sample of SiO2?
____________________
2. How many O atoms is that equivalent to?
____________________
3. How many grams of O atoms would that be?
____________________
4. How many moles of carbon were in your group’s
sample of sucrose?
____________________
5. How many moles of hydrogen were in your group’s
sample of water?
____________________
6. How many atoms of oxygen were in your group’s
sample of sugar?
____________________
7. The ASU Porter-Henderson library has about 1.5 million items (books, journals, manuscripts,
et cetera) in its collection. What fraction of a mole of library items does this represent? If
1.5 million items occupy 3 stories in our library, how many stories would be necessary to
have a mole of library items?
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Data Table 1: One Mole Samples of Various Substances
Substance
Formula
Density
Silica (sand)
SiO2
2.6
g/mL
Sucrose (sugar)
C12H22O11
1.59
g/mL
Sodium chloride
NaCl
2.17
g/mL
Water
H2O
1.00
g/mL
Ethanol
C2H5OH
0.789
g/mL
Zinc, bar
Zn
7.13
g/mL
Zinc, mossy
Zn
Aluminum,
granular
Al
Aluminum, foil ball
Al
Methyl salicylate
(oil of wintergreen)
C8H8O3
1.17
g/mL
Magnesium nitrate
hexahydrate
Mg(NO3)2.6H2O
1.64
g/mL
Copper sulfate
pentahydrate
CuSO4.5H2O
2.28
g/mL
Benzoic acid
C6H5COOH
1.32
g/mL
2.7
g/mL
Molar Mass
Actual Mass
Actual moles
Approx. Volume
No. of Formula
Units in Sample