The Mole Concept

The Mol Concept

Introduction:

The smallest drop of water that the naked eye can see is made up of billions and billions of water molecules. The "mol concept" is a tool that we can use to better grasp such astronomical numbers. A mol is a unit that is used to represent a very large number of atoms or molecules. One mol of any substance is 6.02 x 10²³ (Avogadro's number) particles of that substance. Just as you always assume that there are 12 eggs in a dozen, there will always be 6.02 x10²³ particles in 1 mol of any substance. To give you an idea of how large of a number that really is, if you had a mol of baseballs, it would cover the entire surface of our planet to a height of 100 miles (that's were the space shuttle orbits)!

The molar mass of an element is its atomic weight on the periodic table expressed in grams per mol. For example, the molar mass of carbon is defined as 12.0000 g/mol. The molar mass of a compound is the formula weight in grams for one mol of that substance. Some examples are shown below:

Molecular Formula	Formula Weight	Molar Mass
NaCl	Na Cl (1 x 23.0) + (1 x 35.5) = 58.5 amu	58.5 g/mol
CaCl₂	Ca Cl (1 x 40.0) + (2 x 35.5) = 111 amu	111 g/mol
Na₃(PO₄)₂	Na P O (3 x 23.0) + (2 x 31.0) + (8 x 16.0) = 259 amu	259 g/mol

The molar mass of an element or compound can be used as a conversion factor between grams and mols. For example, how many grams are in 3 mols of CaCl₂?

Or, how many mols of CaCl₂ are in 55.5 g of CaCl₂?

Mol relationships in a chemical reaction can be determined by looking at the balanced reaction equation as shown below for the reaction of aluminum (Al) with hydrochloric acid (HCl) to produce aluminum chloride (AlCl₃) and hydrogen gas (H₂):

A balanced reaction equation has numbers in front of each substance called coefficients. If there is no number in front of a substance, assume the coefficient to be 1. These coefficients tell us the ratio of how many atoms or molecules of each substance will be consumed and produced in that chemical reaction. From the reaction equation above, we can see that for every 2 mols of Al, we will produce 2 mols of AlCl₃. This mol relationship can also be used as a conversion factor. There are two conversion factors that we can derive for each reactant and product in this balanced reaction equation:

mol Al 2 mol AlCl₃	2 mol Al 3 mol H₂	6 mol HCl 2 mol AlCl₃	6 mol HCl 3 mol H₂
2 mol AlCl₃ 2 mol Al	3 mol H₂ 2 mol Al	2 mol AlCl₃ 6 mol HCl	3 mol H₂ 6 mol HCl

The mol to mol relationships or equalities can be used as conversion factors between mols of one substance to mols of another substance in the same chemical reaction. For example, if you started with 1.0 mol of Al in the reaction above, how many mols of H₂ gas would be produced?

Or, if you want to produce 4.0 mols of AlCl₃, how many mols of Al would you need to start with?

Purpose:

In this experiment, you will be conducting two experiments to convince yourself that the molar relationships discussed above do indeed work. In the first experiment, you will be reacting sodium bicarbonate (NaHCO₃) with hydrochloric acid (HCl) to produce sodium chloride, water, and carbon dioxide:

In the second experiment, you will be reacting sodium carbonate (Na₂CO₃) with hydrochloric acid (HCl) to produce sodium chloride, water, and carbon dioxide:

Note that although the products are the same in both reactions, the molar ratios are different. In the first experiment, one mol of NaCl is produced for each mol of NaHCO₃. However, in the second experiment, 2 mols of NaCl are produced for each mol of Na₂CO₃.

At the beginning of the experiment, you will obtain the mass in grams of the sodium bicarbonate or the sodium carbonate. Using the conversion factors discussed above, you will be able to carry out the following conversions:

The grams of NaCl that you determine in these calculations is called the theoretical yield. At the end of the experiment, you will determine the mass in grams of the sodium chloride product and 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, you very rarely obtain the theoretical yield. The percent yield is calculated using the following equation:

Percent yield = (Actual yield/Theoretical yield) x 100%

Procedure:

Thoroughly clean a large Pyrex test tube with soap and water. Then rinse it with 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' your glassware to ensure the removal of all moisture. CAUTION: The glassware will be very HOT - use your tongs when handling!
Allow the test tube to cool to room temperature. Then add a single boiling chip and measure the combined mass to the nearest 0.001 g. (Record your data in a table similar to the one below.)
With a scoopula, add approximately 1.5 g of sodium bicarbonate, NaHCO₃, to the test tube and read the mass to the nearest 0.001 g. Note: Do not try to measure exactly 1.500 g. Your measurement should be about 1.5 g but recored exactly. For example, 1.441 g or 1.582 g are fine.
Obtain about 10 mL of 6M hydrochloric acid in your 10 mL graduated cylinder. CAUTION: HCl causes acid burns - avoid contact with your skin. Use a pasteur (disposable) pipette to slowly add the acid to the NaHCO₃.
When you add the acid to the NaHCO₃, you should observe 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.
Swirl the contents of the test tube to make sure the HCl has come in contact with all of the NaHCO₃. If any unreacted NaHCO₃ remains (bubbles), add a few more drops of HCl to complete the reaction.
Clamp the test tube at a 45° angle and gently heat the liquid over the Bunsen burner until it boils. 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.
Allow the test tube to cool to room temperature and then measure its mass to the nearest 0.001 g.
Reheat the sample strongly for 2-3 minutes. Allow it to cool to room temperature and reweigh. If this weight does not agree to within 0.01 g with the weight in Step 8, reheat and remeasure the mass until two consecutive weights are within 0.01 g of each other. This is known as weighing to constant dryness and 'proves' that all of the water is gone.
Repeat Steps 1-9 using sodium carbonate, Na₂CO₃.

Data Table

a.	Weight of empty Pyrex test tube
b.	Weight of empty Pyrex test tube and boiling chip
c.	Weight of Pyrex test tube, boiling chip, and sample
d.	Mass of sample (c - b)
e.	Weight of Pyrex test tube, boiling chip, and NaCl (after first heating)
f.	Weight of Pyrex test tube, boiling chip, and NaCl (after second heating)
g.	Mass of NaCl (f - b)

Calculations and Conclusions:

Use your initial mass of NaHCO₃ to calculate the number of mols of NaHCO₃ that were used in this reaction.
Use your final mass of NaCl to determine the number of mols of NaCl that were produced.
According to the balanced equation for this reaction, what would you expect the molar ratio of NaHCO₃ to NaCl to be? Do your results agree with this?
What was your percent yield of NaCl? List some possible reasons for your yield to be lower or higher than 100%.
Use your initial mass of Na₂CO₃ to calculate the number of mols of Na₂CO₃ that were used in this reaction.
Use your final mass of NaCl to determine the number of mols of NaCl that were produced.
According to the balanced equation for this reaction, what would you expect the molar ratio of Na₂CO₃ to NaCl to be? Do your results agree with this?
What was your percent yield of NaCl? List some possible reasons for your yield to be lower or higher than 100%.

(Updated 6/7/07 by C.R. Snelling)