CHEM 1030L
Fundamentals of Chemistry Laboratory
The Mole Concept
Background
The Mole Concept
The smallest drop of water that the naked eye can see is made up of billions and billions of water
molecules. The "mole concept" is a tool that is used to better grasp such astronomical numbers. A mole
is a unit that is used to represent a very large number of atoms or molecules. One mole of any
substance is 6.02 x 1023 (Avogadro's number) particles of that substance. Just as there are always 12
eggs in a dozen eggs, there will always be 6.02 x1023 particles in 1 mole of any substance. To give an
idea of how large of a number that really is...if all of the people now alive on the earth started counting
Avogadro's number of peanuts at a rate of two peanuts per second, it would take approximately 2.6
million (2,6000,000) years to count them all. That's a lot of peanuts!
Molar Mass
Avogadro's number has been chosen so that the atomic weight for an element also serves as the molar
mass of that element, expressed in grams per mole. For example, from the periodic table it can be seen
that the atomic weight of carbon is 12.0 amu/atom C and the molar mass of carbon is 12.0 g/mol C.
For a chemical compound, the formula weight (or formula mass) of a substance can be obtained by
adding up the atomic masses for all elements in the substance. Very conveniently, the molar mass of a
compound is equal to the formula weight of the compound, but expressed in grams per mole. Some
examples are shown in the following table.
Chemical
Formula
NaCl
CaCl2
Na3PO4
Formula Weight
Na
Cl
(1 x 23.0) + (1 x 35.5) = 58.5 amu
Ca
Cl
(1 x 40.0) + (2 x 35.5) = 111.0 amu
Na
P
O
(3 x 23.0) + (1 x 31.0) + (4 x 16.0) = 164.0 amu
Molar Mass
58.5 g/mol
111.0 g/mol
164.0 g/mol
Using Molar Mass as a Conversion Factor
The molar mass of an element or compound can be used as a conversion factor between grams and
moles of that substance. For example, the molar mass of NaCl can be written either of the following
ways.
58.5 g NaCl
1 mole NaCl
OR
1 mole NaCl
58.5 g NaCl
The version of the conversion factor used in a calculation depends on the units that should cancel out
during the calculation, as shown in the following examples.
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Example Problem 1, convert 1.453 g NaCl to moles NaCl.
1.453 g NaCl x
1 mole NaCl
58.5 g NaCl
=
0.0248 mol NaCl
In answering Problem 1, notice that g NaCl needed to cancel out. Therefore, the conversion factor
needed to have g NaCl on the bottom. The units that did not cancel out were mol NaCl, which were the
units required to answer the question.
Example Problem 2, convert 1.45 moles of NaCl to grams.
1.45 mol NaCl x
58.5 g NaCl
1 mol NaCl
=
84.8 g NaCl
In answering Problem 2, the units of mol NaCl needed to cancel out. Therefore, the conversion factor
needed to have mol NaCl on the bottom. The units that did not cancel out were g NaCl, which were the
units required to answer the question.
Hint: When working a mole conversion problem, do not start with a conversion
factor. The chances are 50/50 that the conversion factor will be upside down.
Instead, start working the problem with a measured quantity, such as grams or
moles.
Today’s Experiment
In this experiment, students will work in pairs to react hydrochloric acid (HCl aq) with one of two
compounds - either sodium carbonate (Na2CO3) or sodium hydrogen carbonate (also known as sodium
bicarbonate - NaHCO3) to produce sodium chloride, water and carbon dioxide.
Na2CO3 (s) + 2 HCl (aq) —> 2 NaCl (aq) + CO2 (g) + H2O (l)
OR
NaHCO3 (s) + HCl (aq) —> NaCl (aq) + CO2 (g) + H2O (l)
To begin the experiment, each student will obtain the mass of the sodium carbonate or sodium
bicarbonate. The conversion factors discussed above will be used to convert that mass of each reactant
(Na2CO3 and NaHCO3) to the mass of NaCl that theoretically should be produced by the reaction. This
is called the theoretical yield. After the reaction is complete, the mass of NaCl actually produced by
the reaction will be determined by weighing the product. This is called the actual yield.
According to the law of conservation of mass, the actual yield should be equal to the theoretical yield.
However, due to human and experimental errors, it very seldom does. Percent yield is a measure of
how close the actual yield is to the theoretical yield and is calculated using the following equation.
Percent yield =
Revised 10/17/2015 – T. Ekman
___actual yield___ x
theoretical yield
100
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Procedure
Set up apparatus as described below to match the example apparatus on the instructor's bench.
1. Look at the underside of the student lab bench to find the rectangular metal bin that holds the
ring stand pole. Slide the pole out of the opening at the end of the bin.
2. Insert the beveled end of the pole into the metal socket on the lab bench top.
3. Remove the following equipment from the drawer.
 Bunsen burner with rubber hose
 striker
 two-jaw clamp
IMPORTANT: Use a clamp that does not have a plastic coating on the jaws. The
plastic will melt and stick to the test tube.
 large Pyrex test tube
 250-mL beaker
4. Attach the free end of the rubber hose to the gas tap and place the Bunsen burner in front of the
ring stand pole.
5. Place the clamp on the ring stand pole. Clamp the test tube at the open end and tilt it at a 45°
angle. Rotate the test tube that the open end is not pointed at yourself or the lab partner.
6. Use a ruler to adjust the height of the clamp so that the bottom of the test tube is 8-10 cm above
the top of the Bunsen burner.
7. Thoroughly clean the test tube with soap and water and rinse with tap water and then distilled
water. Fold a paper towel into a long thin strip and use this to remove as much of the water as
possible. Finally, use a Bunsen burner to 'flame dry' the test tube to ensure the removal of all
moisture. CAUTION: The glassware will be very HOT – do not attempt to remove it from the
clamp until it is cool enough to handle.
8. Allow the test tube to cool to room temperature. Remove it from the clamp and rest it in the
250-mL beaker.
9. Add a single boiling chip to the test tube. Weigh the test tube and boiling chip to the nearest
0.001 g.
Remember that the easiest way to weigh a test tube is to first zero the beaker on
the balance and then place the test tube in the beaker. The test tube cannot touch
the inside of the balance compartment.
10. As directed by the instructor, use a scoopula to add approximately 1.5 g of either Na2CO3 or
NaHCO3 to the test tube. Zero the beaker and read the mass of the test tube, boiling chip, and
reactant powder to the nearest 0.001 g. Note: Do not try to measure exactly 1.500 g. The mass
does not have to be exactly 1.500 g, but it is important to know exactly what the mass actually
is. For example, 1.442 g or 1.587 g are acceptable masses – as long as they are accurately
recorded.
11. Use a 10-mL graduated cylinder to obtain about 8 mL of 6M hydrochloric acid.
CAUTION: HCl causes acid burns - avoid skin contact.
12. Place the beaker and test tube on the bench and lower the snorkel hood over them. Use a
disposable pipette to slowly add the acid to the Na2CO3. Feel the end of the test tube to observe
whether the reaction is exothermic or endothermic. Also observe the temperature of the lab
partner's test tube. Addition of the HCl will cause the evolution of gas (bubbles). Continue to
add the acid slowly until the reaction is complete (bubbling has stopped). Do not add more
acid than is needed. A white solid may form as the reaction nears completion.
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13. Swirl the contents of the test tube to make sure the HCl has reacted with all of the solid. New
bubbles will indicate that unreacted solid is present. If new bubbles are observed, add a few
more drops of HCl to complete the reaction.
14. Clamp the test tube near the opening and at a 45° angle. Gently heat the liquid over the Bunsen
burner until it boils. Initially, use a small flame approximately 2-3 inches high. Take care to
avoid loss of liquid from boiling over. If the liquid begins to splatter, remove the heat
immediately. Lower the flame and then continue to heat. Continue to dry the solid slowly until
all moisture appears to have evaporated.
15. Allow the test tube to cool to room temperature and measure its mass to the nearest 0.001 g.
16. Reheat the sample strongly for 2-3 minutes, but do not allow the Bunsen burner to heat one
part of the test tube for very long, which can cause melting or shattering. If demonstrated by
the lab instructor, use the "flame brushing" technique.
17. Allow the test tube to cool to room temperature and re-weigh it.
18. If the second mass does not agree to within 0.01 g with the first mass, reheat and remeasure the
mass until two consecutive masses are within 0.01 g of each other. This is known as heating to
constant mass and "proves" that all of the water has been evaporated.
19. Before leaving lab, obtain data from the lab partner for the other solid that was reacted.
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