Fundamentals of Chemistry 1030
The Mole Concept
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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 we can use 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 you would always assume that there are 12 eggs in a dozen, there will always be 6.02 x1023 particles in 1 mole of any substance. To give you 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. That's a lot of peanuts!

Molar Mass

The molar mass of an element is its atomic weight on the periodic table expressed in grams per mole. For example, the molar mass of carbon is 12.0 g/mol. The molar mass of a compound is the formula weight in grams for one mole 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
Na3PO4      Na              P                 O 
(3 x 23.0) + (1 x 31.0) + (4 x 16.0) = 164 amu
164 g/mol
 

Using the 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. For example, how many grams are in 3 moles of CaCl2?


 
 
 
 

Or, how many moles of CaCl2 is 55.5 g of CaCl2?


 
 

Mole Relationships in a Chemical Reaction

Mole 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 elements 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 moles of Al, we will produce 2 moles of AlCl3. This mole relationship can also be used as a conversion factor. There are several conversion factors that we can derive from this balanced reaction equation:

Using the Mole Relationship as a Conversion Factor

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


 
 

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


 
 
 
 

Today’s Experiment

In this experiment, you will be reacting sodium bicarbonate (NaHCO3) with hydrochloric acid (HCl) as shown below to produce sodium chloride, water and carbon dioxide:

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

 The grams of NaCl that you determine in this calculation 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, it very seldom is. The percent yield is calculated using the following equation:

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

Laboratory Procedures


1. 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!

2.  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 1.50 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 1.500 g but recorded precisely.  For example, 1.447 g or 1.528 g are correct ways (mass has been recorded to 0.001g.

3. Obtain about 5 mL of 6M hydrochloric acid in your 10 mL graduated cylinder. CAUTION:  HCl (aq) 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.

4. 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.

5. Clamp the test tube near its 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.

6. Allow the test tube to cool to room temperature and then measure its mass to the nearest 0.01 g.

7 .Reheat the sample strongly for 2-3 minutes. Allow it to cool to room temperature and re-weigh.
 
8. If this weight does not agree to within 0.01 g with the weight in Step 7, 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.                                        Data Table  
a. Mass of  test tube and boiling chip  
b. Mass of test tube and boiling chip and NaHCO3  
c. Mass of NaHCO3 (b - a)  
d. Mass of  test tube and boiling chip NaCl (after first heating)  
e. Mass of  test tube and boiling chip NaCl (after second heating)  
f. Mass of NaCl (e - a)  

SHOW ALL CALCULATIONS IN LAB NOTEBOOK.

Issues to be addressed in your conclusion...

Use your initial mass of NaHCO3 to calculate the number of moles of NaHCO3 that were used in this reaction.

Use your final mass of NaCl to determine the number of moles 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%.