Experiment 11   Preparation of a Compound and Calculation of an Empirical Formula

 

 

PRIOR READING

Be sure to read and understand the material in chapter 3, section 3.2. 

 

 

INTRODUCTION

The law of definite proportions states that a compound always contains two or more elements combined in the same definite proportion by mass.  Compounds that are composed of molecules are called covalent compounds; compounds that are composed of ions are called ionic compounds.  Two examples are water and magnesium oxide.  Both are represented by their chemical formulas H2O, and MgO, respectively.  A chemical formula indicates the relative number of atoms in a compound and is termed empirical formula.  Water contains 2 hydrogens per one oxygen and magnesium oxide contains 1 magnesium per 1 oxygen.  The chemical formula that indicates the actual number and type of atoms in a compound is called molecular formula. 

 

Consider for example a compound composed of two atoms, H and O.  We know that water consists of two atoms of hydrogen and one atom of water, giving it the molecular formula of H2O.  Another compound composed of these same elements, but in different proportions, is hydrogen peroxide.  The molecular formula for hydrogen peroxide is H2O2, saying that there are actually 2 hydrogens and 2 oxygen atoms bonded to give H2O2.  The empirical formula for this compound is HO, showing the smallest whole-number ratio of atoms in a compound. 

 

The empirical formula of a compound can be determined in the laboratory by its synthesis from the elements by measuring the mass of each component in a compound.  This is used to calculate the number of moles of each element, and finally the empirical formula of the compound.  The empirical formula of a compound specifies the simplest, whole-numbered ratio of elements.  Percent composition of a compound is the mass in grams of one element contained in 100 g of the whole compound. 

 

In this experiment, magnesium is reacted with oxygen from the air to form magnesium oxide, according to:

                                    2 Mg  +  O2    2 MgO

 

In today’s experiment, the empirical formula of magnesium oxide can be determined from the mass of the magnesium initially present and the mass of the magnesium oxide compound. 

 

We will also calculate the percent error for the magnesium in the sample.  The percent error will show us how far off the experimental value is from the predicted, theoretical, value.  The percent error is calculated using the following equation:

 

            Percent error = Theoretical value - Experimental value   x 100

                                                            Theoretical value

 

 

 

MATERIALS

Crucible and cover

Magnesium; needs to be polished

 

 

EXPERIMENTAL PROCEDURE

Record all data and observations directly into your notebook. 

 

1.  Heat a crucible by placing it on a clay triangle over a Bunsen burner.  Allow it to cool to room temperature.  Weigh the crucible to the nearest 0.0001 g and record the weight into your notebook.  Handle the crucible with tongs only. 

 

2.  Add about 0.20 g of polished magnesium (use steel wool and be careful not to scratch the counter tops) and weigh again accurately to the nearest 0.0001 g and record the weight into your notebook. 

 

3.  Again, heat the crucible on a clay triangle using your Bunsen burner.  Carefully place the lid on the crucible.  Heat slowly, occasionally lifting the lid to allow air to react with the magnesium. 

Try to avoid the loss of magnesium oxide smoke.  Continue heating until no change is apparent in the magnesium ash (about five to ten minutes).  Allow the crucible to cool to room temperature.  Add a few drops of water * and reheat the crucible, slowly first, then strongly for about five more minutes. 

 

4.  Allow the crucible to cool to room temperature and record its mass to the nearest 0.0001 g.  Reheat to ensure completion. 

 

5.  Complete a second trial following steps 1 through 4.  (You can perform both trials at the same time)

 

*  Because the air contains a large amount of nitrogen gas, a portion of the magnesium metal will react with N2 to form magnesium nitride.  Magnesium nitride is decomposed by its reaction with water to form magnesium oxide and ammonia. 


 

WASTE DISPOSAL

Dispose of the magnesium oxide in the container designated. 

 

 

 

Prelaboratory Questions   Empirical Formula

1.  An unknown oxide of mercury decomposes when heated to form mercury metal and oxygen gas (O2).  When a 1.048 g sample of this unknown is heated, 0.971 g of mercury remains. 

a.  Calculate the moles of mercury and moles of oxygen in the compound. 

b.  What is the empirical formula of this oxide? 

2.  a) A side reaction in today’s experiment occurs between some of the magnesium and the nitrogen gas, N2. Write a balanced chemical equation for this reaction.

b) The magnesium nitride can be converted to magnesium oxide by the addition of water.  Write a balanced chemical equation for this reaction. 

3.  Suppose 1.087g of magnesium is heated in air.  What is the theoretical amount of magnesium oxide that should be produced?

4.  Calculate the percent by mass of molybdenum and sulfur in Mo2S5.

 

 

 

 

Data and Analysis:   Empirical Formula

 

Data Collection

You and you partner are to perform 2 trials.  Prepare a table in your notebook to include all pertinent information.  Calculate the % by mass of Mg in the oxide and compare it to the % by mass that you have observed.  Then calculate % by mass of oxygen in the oxide.  What is the empirical formula for you oxide?  What is your percent error?             

 

Questions

1.  If the magnesium is burned uncontrolled (burns brightly), how will this error affect the reported mass of the

a)  magnesium in the sample?

b)  oxygen in the sample?

c)  magnesium oxide?

Please explain. 

2.  What error would result if the water were not removed completely in the last heating?

3.  What error would occur if magnesium nitride were not decomposed in step 3?