MOLE RELATIONSHIP IN A CHEMICAL REACTION

            The Law of Conservation of Matter states that “matter is neither created nor destroyed in a chemical reaction.”  In other words, the total mass of the reactants must equal the total mass of the products in a chemical reaction.  Chemical equations are balanced so that they do not contradict the law of conservation of matter.  The coefficients used to balance an equation also give the relative number of moles of reactants and products.

 

            In this activity, you will test the Law of Conservation of Matter by causing a chemical reaction to occur with a given amount of reactant.  You will then carefully determine the mass of one of the products.  With these measurements, you will be able to calculate the moles of one reactant and one product and compare the number of moles.  From the balanced equation, you should be able to see the relationship between the number of moles of a reactant and the number of moles of a product.

 

OBJECTIVES

- predict a balanced equation for the reaction taking place

- react a known mass of Na2CO3 with excess HCl

- calculate the mole ratio between Na2CO3 and NaCl

- determine whether your results support the Law of Conservation of Matter

 

EQUIPMENT

- goggles & apron                      - balance                       - forceps or tongs

- evaporating dish                      - lab burner                   - watch glass

- dropper pipet                           - 250 mL beaker

 

PROCEDURE

* SAFETY GOGGLES AND LAB APRON MUST BE WORN AT ALL TIMES DURING THIS EXPERIMENT! *

 

1.  Clean and dry an evaporating dish.  Determine the mass of the empty, dry evaporating dish to the nearest
     0.01 g.

2.  With a spatula, add about 1.5 grams of sodium carbonate to the evaporating dish, and read the mass to the
     nearest 0.01 g.  (NOTE:  You should not attempt to measure exactly 1.50 g since it is only a reference point. 
     For example, mass readings of 1.72 and 1.38 would be equally acceptable.)

3.  Cover the evaporating dish with a watch glass.  Using the dropper bottle, carefully add hydrochloric acid to
     the evaporating dish (that already has the Na2CO3 in it).

CAUTION:  HCl causes burns;  avoid skin & eye contact.  Rinse spills with plenty of water. 
      Allow the drops to enter the lip of the evaporating dish so that they gradually flow down the side.

4.  Continue adding the acid slowly until the reaction has stopped.  Do not add more acid than is needed to
     complete the reaction.  (If you do add more than is needed, the rest of the lab will take longer.)

5.  Tilt the dish from side to side to make sure the HCl has reacted all of the Na2CO3.  If any unreacted Na2CO3
     remains, add a few more drops of HCl to complete the reaction.  Remove the watch glass cover.  Rinse the
     underside of the watch glass with a very small amount of water.  Be careful to wash all material into the
     evaporating dish.

6.  Heat the liquid in the evaporating dish with a low flame on the lab burner until it boils GENTLY.  Take care to
     avoid loss of liquid from boiling over.  Continue to dry the solid slowly until all moisture appears to have
     evaporated.

7.  Remove the dish from the heat and allow it to cool.  Then measure and record the mass to the nearest
     0.01 g.

8.  After massing, the contents of the dish may be rinsed down the drain using plenty of water.  Clean all lab
     equipment and return it to the container.

 

DATA TABLE

Mass of empty evaporating dish

____________________g 

Mass of evaporating dish & Na2CO3

____________________g 

Mass of Na2CO3

____________________g 

Mass of evaporating dish & NaCl

____________________g 

Mass of NaCl

____________________g 

 

QUESTIONS AND CALCULATIONS

1.  One of the products in this reaction was NaCl, the other two products were gases.  These gases are
     also produced in a combustion reaction.  Write the balanced equation for the reaction in this
     experiment.

 

 

2.  From your balanced equation, what is the mole ratio between Na2CO3 and NaCl?

 

3.  Suppose you had started with 3.25 moles of sodium carbonate.  How many moles of sodium
     chloride would you expect to be formed?  If you started with X moles of Na2CO3, how many moles
     of NaCl would you expect to be formed?  Explain.

 

 

 

4.  Calculate the number of moles of Na2CO3 used in this reaction.

 

 

 

5.  Calculate the number of moles of NaCl produced in this reaction.

 

 

 

6.  From the data you collected in the lab, what is the mole ratio between Na2CO3 and NaCl?

 


7.  How does the mole ratio from the balanced equation compare to the mole ratio from the
     experiment?

 

 

8.  Using the mass of Na2CO3 that you actually used in the experiment, determine the theoretical yield
     of NaCl in this experiment.  (You will need to do a stoichiometry problem starting with your mass of Na2CO3
     and determine the number of grams of NaCl that should have been produced.)

 

 

 

 

9.  Compare the theoretical yield with your actual yield.  What is your percent yield?



 

10.  Was your percent yield more or less than 100 %?  Explain what your percent yield means.
(Explain why – in terms of the lab procedure - your percent yield was less than or more than 100 %.)

 

11.  Write a paragraph describing the observations of this chemical reaction.  Also include in this
       paragraph:
            - What are the indicators that a chemical reaction occurred?

- What were two (2) sources of error in this experiment?  (Assume balances are correct and
             that all instrument readings were done correctly.)

- What could be done to prevent these errors if you did the experiment again?

 

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