EXPERIMENT
Molarity: Conversions and
Mass Determination
Hands-On Labs, Inc. Version 42-0206-00-02
Review the safety materials and wear goggles when
working with chemicals. Read the entire exercise
before you begin. Take time to organize the materials
you will need and set aside a safe work space in
which to complete the exercise.
Experiment Summary:
You will calculate the number of moles and atoms
of a compound from the mass of the sample and
the molar mass. You will also dry a sample of
hydrated alum to calculate the number of moles of
water present in the sample. You will determine the
empirical formula of the hydrate.
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Experiment
Molarity: Conversions and Mass Determination
Learning Objectives
Upon completion of this laboratory, you be able to:
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Define the concept of a mole.
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Distinguish between the terms: atomic mass, molecular mass, and molar mass.
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Convert between mass, moles, and atoms.
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Define and name hydrates by the number of moles of water held in a compound.
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Calculate the number of moles and atoms of a substance.
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Calculate the number of moles of water released by a hydrate.
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Determine the empirical formula of the hydrate from the formula of the anhydrous compound
and the experimental data.
Time Allocation: 2.5 hours
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Experiment
Molarity: Conversions and Mass Determination
Materials
Student Supplied Materials
Quantity
1
1
1
Item Description
Aluminum pie pan
Box of matches
Cotton ball
HOL Supplied Materials
Quantity
1
2
1
1
1
1
1
1
1
Item Description
Alum crystals, KAI(SO4)2 - 5 g
Aluminum cups, 2 oz
Burner fuel
Burner stand
Digital scale, precision
Glass stir rod
Pair of gloves
Pair of safety goggles
Test tube clamp
Note: To fully and accurately complete all lab exercises, you will need access to:
1. A computer to upload digital camera images.
2. Basic photo editing software such as Microsoft Word® or PowerPoint®, to add labels, leader
lines, or text to digital photos. 3. Subject-specific textbook or appropriate reference resources from lecture content or other
suggested resources. Note: The packaging and/or materials in this LabPaq kit may differ slightly from that which is listed
above. For an exact listing of materials, refer to the Contents List included in your LabPaq kit.
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Experiment
Molarity: Conversions and Mass Determination
Background
The Mole
A single atom or a single molecule is so small that chemists are seldom able to work with one
at a time. Even when weighing the smallest quantities of substances, numerous atoms and
molecules are present. A unit of measure called a mole is used to combat this problem, allowing
for successful work with defined quantities of atoms and molecules.
A mole (η) is a unit of measure, describing the amount of a chemical substance that contains
as many atoms, molecules, and ions as there are in exactly 12 grams of pure carbon (12C). This
amount of atoms (6.022 × 1023) is referred to as Avogadro’s number. One mole of a substance
(element or compound) is equal to its atomic mass (molecular mass), expressed in grams. The
atomic masses shown on the periodic table are the average masses of the various isotopes of
each element, expressed in atomic mass units, or amu.
One atomic mass unit is equivalent to 1 gram per mol. Thus, the atomic mass units shown on
a periodic table are equal to the molar masses of the elements or the mass in grams of each
element that contains one mole of atoms of that element. For example, the element nitrogen
has a molar mass of 14.007 grams, thus 1 mole of nitrogen is equal to 14.007 grams. Likewise,
the compound H2O has a molar mass of 18.015 (H + H + O = 1.008 + 1.008 + 15.999), thus 1 mole
of H2O is equal to 18.015 grams. The molar mass (atomic mass) of each element is found in the
periodic table. See Figure 1.
Figure 1. Periodic Table of Elements. The atomic mass (atomic weight) of an element is equal
to the mass in grams required to equal 1 mole of the substance. Click to Download Printable
Version.
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Experiment
Molarity: Conversions and Mass Determination
Conversions
Molar mass conversions are common calculations in chemistry laboratories. Example calculations
are provided below to further explore the relationships between molar masses, moles, grams,
and atoms. For each example, assume that a vial contains 12.5 grams of magnesium chloride
(MgCl2), and you need to know how many moles and atoms of magnesium and chlorine are in
the vial.
Example #1: Calculating the Molar Mass of a Compound
What is the molar mass of MgCl2?
1. Identify the atomic mass of each element in the compound using the periodic table.
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For MgCl2:
2. Calculate the molar mass of each element by multiplying the atomic mass by 1 g/mol.
3. Add the molar mass of each element in the compound together.
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In the compound MgCl2 there is 1 mole of magnesium and 2 moles of chlorine:
The molar mass of MgCl2 is equal to 95.211 g/mol.
Example #2: Calculating the Number of Moles of a Compound in a Sample
How many moles of MgCl2 are present in a sample weighing 12.5 g?
1. Calculate the molar mass of the compound using the periodic table.
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From Example #1, the molar mass of MgCl2 is 95.211 g/mol.
2. Use the molar mass to convert the mass (grams) of the sample to moles.
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The mass of the MgCl2 sample in this example is 12.5 g.
There are 0.131 moles of MgCl2 in 12.5 g of MgCl2.
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Experiment
Molarity: Conversions and Mass Determination
Example #3: Calculating the Number of Moles of an Element in a Sample
How many moles of Mg are present in 12.5 g MgCl2? How many moles of Cl are present?
1. Calculate the molar mass of the compound using the periodic table.
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From Example #1, the molar mass of MgCl2 is 95.211 g/mol.
2. Use the molar mass to convert the mass (grams) of the sample to moles.
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From Example #2, there are 0.131 mol MgCl2 in a 12.5 g sample.
3. Calculate how many moles of each element are present in the moles of the compound.
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For MgCl2, calculate the number of moles of Mg present in 0.131 moles of MgCl2:
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Calculate the number of moles of Cl present in 0.131 moles of MgCl2:
There are 0.131 moles of Mg and 0.262 moles of Cl in 12.5 g of MgCl2.
Example #4: Calculating the Number of Atoms of an Element in a Sample
How many atoms of Mg are present in 12.5 g MgCl2? How many atoms of Cl are present?
1. Calculate the molar mass of the compound using the periodic table.
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From Example #1, the molar mass of MgCl2 is 95.211 g/mol.
2. Use the molar mass to convert the mass (grams) of the sample to moles.
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From Example #2, there are 0.131 mol MgCl2 in a 12.5 g sample.
3. Calculate how many moles of each element are present in the moles of the compound.
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From Example #3, there are 0.131 mol Mg and 0.262 mol Cl in 0.131 mol MgCl2.
4. Apply Avagadro’s number to convert the moles of each element to numbers of atoms.
There are 7.89 x 1022 atoms of Mg and 1.578 x 1023 atoms of Cl present in 12.5 g of MgCl2.
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Experiment
Molarity: Conversions and Mass Determination
Another Example of Converting Grams to Moles
Potassium aluminum sulfate (KAI(SO4)2), commonly referred to as alum, has a wide range of uses
in industry, cosmetics, cooking, and even taxidermy. How many moles of potassium aluminum
sulfate are in a jar containing 10 g of the substance?
First, determine the molar mass of the compound.
Next, use the molar mass to convert the grams to moles.
The mole is attributed to an 18th century
Italian scientist whose full name, Lorenzo
Romano Amedeo Carlo Avogadro di Queregna e
di Cerreto, fits the long number that now bears his
name. Avogadro’s findings were ignored for nearly
100 years before the French chemist and Nobel
laureate, Jean Baptiste Perrin, proposed that
the quantity of molecules be called “Avogadro’s
Constant.”
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Experiment
Molarity: Conversions and Mass Determination
Hydrates
In the experiment, the mole concept is used to determine how many moles of water are present
in a sample of hydrated potassium aluminum sulfate. A hydrate is a compound in which water
molecules are a part of the crystalline structure. The water is loosely held in the compound and
is easily separated from the compound upon heating. It is important to note that a hydrate is
different than a sample that is merely wet. Hydrates are named by adding the Greek prefixes
mono-, di-, tri-, penta-, hexa-, hepta-, etc., to the end of the standard name of the compound,
followed by the word “hydrate,” to describe moles of water held in the compound. For example:
The formulas indicate that 1 mole of CuSO4•5H2O contains 1 mole of CuSO4 and 5 moles of H2O;
and 1 mole of MgSO4•7H2O contains 1 mole of MgSO4 and 7 moles of H2O. If we were to heat
these two hydrates, the water molecules would be released:
As shown in the equations, for each mole of hydrated copper (II) sulfate that is heated, 5 moles of
water are released, leaving behind 1 mole of copper (II) sulfate. Likewise, for each mole of hydrated
magnesium sulfate that is heated, 7 moles of water are released, leaving behind 1 mole of
magnesium sulfate.
In this experiment, the mass of three items will be used to calculate the number of moles and
atoms of each element present in the chemical compound. In addition, the chemical formula of a
hydrated potassium aluminum sulfate (alum) will be determined. By calculating the mass of water
released by the hydrated potassium aluminum sulfate after heating and converting the mass of
water to moles of water, the chemical formula of the hydrated compound can be determined. In
this experiment the hydrated potassium aluminum sulfate (alum) will be heated multiple times,
to ensure that all of the water contained in the sample is released. When the sample is heated
and all of the water has been released, the mass of the sample will not change and will remain
constant no matter how many additional times the sample is heated.
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Experiment
Molarity: Conversions and Mass Determination
Exercise 1: Using Mass to Calculate Moles and Atoms
In this exercise, you will determine the mass of three items. You will use the mass and molecular
formula of each item to calculate the moles and atoms of each element present in the item.
Procedure
1. Gather the digital scale, 2 oz. aluminum cup, glass stirring rod, and cotton ball.
2. Use the periodic table in Figure 1 of the Background to calculate the molar mass of the
aluminum cup. Click here to download a printable copy of the periodic table.
Note: The formula of the aluminum cup is provided in Data Table 1.
3. Record the molar mass in Data Table 1 of your Lab Report Assistant.
4. Turn on the digital scale and ensure that it displays 0.00 g.
5. Place the aluminum cup on the scale and determine its mass. Record the mass of the aluminum
cup in Data Table 1.
6. Calculate each of the following and record your findings in Data Table 1.
a. Number of moles of Al2O3 present in the aluminum cup.
b. Number of moles of Al and O present in the aluminum cup.
c. Number of atoms of Al and O present in the aluminum cup.
7. Repeat steps 2 through 6 for the glass stirring rod and the cotton ball in Data Table 1.
Cleanup:
8. Return all items to the kit for future use.
Questions
A. Describe the relationship between moles and atoms.
B. A sample of 2 tsp of sugar (C12H22O11) weighs 9.00 g. Record each step needed to calculate the
moles and atoms of all elements present in the sample. Then, calculate the moles and atoms
of each element in the sample of sugar. Show all work.
C. Which item in Data Table 1 contains the largest quantity of moles?
D. Which item in Data Table 1 contains the least amount oxygen atoms?
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Experiment
Molarity: Conversions and Mass Determination
Exercise 2: Water of Hydration
In this exercise, you will determine the chemical formula of a hydrated potassium aluminum
sulfate sample (alum).
Procedure
1. Put on the safety goggles and gloves.
2. Turn on and tare the digital scale so that it reads 0.00 g.
3. Place the aluminum cup on the scale to obtain the mass and record the mass in Data Table 2
of your Lab Report Assistant.
4. Using the second aluminum cup as a weigh boat, place the aluminum cup on the scale, tare
the scale and measure 2.0 grams of KAl(SO4)2.
5. Transfer the alum to the first aluminum cup and place the aluminum cup containing the
KAl(SO4)2 on the scale to obtain the mass. Record the mass in Data Table 2.
6. Assemble the burner setup and light the fuel, as shown in Figure 2.
a. Gather the burner fuel, aluminum pie plate, burner stand, 2 oz. aluminum cup, and
lighter or matches.
b. Place an aluminum pie plate on a solid work surface away from flammable objects.
c. Set the burner stand towards the back of the pie plate.
d. Use matches or a lighter to ignite the fuel. BE CAREFUL- the flame may be nearly
invisible.
e. Gently slide the fuel under the stand.
f. The small, 2 oz. aluminum cup will be placed over the fuel to extinguish the flame. Set
the aluminum cup next to the burner setup so you are ready to extinguish the flame
at any point.
Figure 2. Burner fuel setup.
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Experiment
Molarity: Conversions and Mass Determination
7. Pick up the aluminum cup containing the KAl(SO4)2 with the test tube clamp, as shown in
Figure 3, and place it onto the burner stand, as shown in Figure 4.
Figure 3. Holding aluminum cup with test tube holder.
Figure 4. Aluminum cup on burner stand with heat source.
Note: You will see the water being released from the alum as it bubbles and evaporates off of the
sample.
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Experiment
Molarity: Conversions and Mass Determination
8. Allow the aluminum cup to heat on the burner for approximately 5 minutes or until all bubbling
has ceased and the sample appears dry. See Figure 5.
Figure 5. Dry sample (KAl(SO4)2) following heating.
9. After the aluminum cup cools for a few minutes, use the test tube clamp to transfer it to the
tared digital scale. Record the mass of the aluminum cup and KAl(SO4)2 in Data Table 2.
10.Repeat steps 7 through 9 another time, to complete the 2nd heating in Data Table 2.
Note: Step 10 is reheating the same sample, NOT starting with a fresh sample.
11. Use the other empty 2 oz. aluminum cup to extinguish the burner fuel flame. See Figure 6.
a. Do not touch the metal stand; it may be hot.
b. Carefully slide the burner fuel canister out from underneath the burner stand. The
sides of the burner fuel canister will be warm, but not hot.
c. Place the aluminum cup directly over the flame to smother it. The cup should rest on
top of the fuel canister, with little or no smoke escaping. Do not disturb the burner
stand and beaker; allow everything to cool completely.
d. Once all equipment is completely cool, remove the aluminum cup and place the
plastic cap back on the fuel. Ensure that the plastic cap “snaps” into place to prevent
fuel leakage and evaporation. The aluminum cup, fuel, and all other equipment may
be used in future experiments.
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Experiment
Molarity: Conversions and Mass Determination
Figure 6. Extinguishing burner with flame.
12.Calculate the mass of water released when the sample was heated by subtracting the mass
of the 2nd heating from the mass of the Aluminum Cup + 2.0 grams of KAl(SO4)2 and record in
Data Table 2.
13.Calculate the molecular mass of H2O and record in Data Table 2.
14.Calculate the moles of water released by heating the sample and record in Data Table 2.
Cleanup:
15.Properly dispose of the cup containing the sample and remaining chemicals.
16.Clean up and return all other items to the kit for future use.
Questions
A. Calculate the moles of anhydrous (dry) KAl(SO4)2 that were present in the sample. Show all
work including units.
B. Calculate the ratio of moles of H2O to moles of anhydrous KAl(SO4)2. Show all work including
units. Note: Report the ratio to the closest whole number.
C. Write the empirical formula for the hydrated KAl(SO4)2, based on your experimental results
and answer to Question 2. Show all work including units. Hint: if the ratio of moles of H2O to
moles of anhydrous KAl(SO4)2 was 4, then the empirical formula would be: KAl(SO4)2•4H2O.
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Experiment
Molarity: Conversions and Mass Determination
D. Describe any visual differences between the hydrated sample and the dried, anhydrous form.
E. How would the following errors affect the empirical formula for the compound?
a. The student ran out of time and did not do the second heating. Explain how this error
will affect the calculation for the number of moles of water in the hydrate? Will the
final answer be artificially high or low? How do you know?
b. The student recorded the mass of the cup + sample incorrectly and started with 2.20
g of hydrated compound but used 2.00 g in the calculations. Explain how this error
will affect the calculation for the number of moles of water in the hydrate? Will the
final answer be artificially high or low? How do you know?
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