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8: LAB 8 - EMPIRICAL FORMULAS AND HYDRATES

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    PURPOSE

    The purpose of this experiment is to:

    • Determine the empirical formula of a hydrate.
    • Calculate the percentage of water in a hydrate.

    INTRODUCTION

    An empirical formula is the smallest whole-number ratio of elements in a compound. It can be determined experimentally by examining a substance's molar ratio of individual components. Consider the following example:

    Example \(\PageIndex{1}\)

    A compound was shown to consist of 0.750 g of carbon and 2.000 g of oxygen. What is the empirical formula?

    Solution

    To calculate the empirical formula, we would first need to convert grams of each element to moles, as shown below:

    0.750 g of C times 1mol of C over 12 g of C is equal to 0.0625 mol of C

    2.00 g of O times 1mol of O over 16 g of O is equal to 0.125 mol of O

    Since we know the ratio of moles, we can automatically determine the ratio of atoms, as 1 mole equals 6.02 × 1023 atoms. Therefore, the ratio of moles of the carbon oxide compound is:

    0.125 mol over 0.0625 mol is equal to two

    Since there are twice as many moles of oxygen as carbon, the empirical formula is CO2.

    In this experiment, you will determine the empirical formula of a hydrate, a compound containing water molecules. First, you will heat the hydrate to drive off the water molecules and then determine the mass of water present in the hydrate by subtracting the mass of the anhydrous compound from the mass of the original hydrate. The empirical formula will be found by comparing the moles of anhydrous compound to the moles of water, as shown in the example below:

    Example \(\PageIndex{1}\)

    2.500 g of a copper(II) chloride hydrate was heated, resulting in 1.972 g of anhydrous compound. Using this information, determine the empirical formula of the copper chloride hydrate.

    Solution

    First, we can find the mass of water driven off:

    2.500 g hydrate – 1.972 g anhydrous compound 0.528 g water

    Next, we can find the molar ratio of the anhydrous compound (AC) to water. Since the anhydrous compound has the formula CuCl2, we can convert grams to moles, as follows:

    1.972 g of AC times 1 mol AC over 134.45 g AC is equal to 0.0147 mol AC

    Similarly, we can find the moles of water:

    0.528 g water times 1 mol of water / 18 g water is equal to 0.0293 mol water

    Comparing the molar ratios gives:

    0.0293 mol over 0.0147 mol is equal to 2

    The empirical formula would therefore be written as CuCl2· 2H2O, indicating that there are two water molecules for every molecule of the anhydrous compound. Since the molar mass of the hydrate is 170.45 g/mol (adding up the molar masses from Cu, 2 Cl, and 2 H2O), the percentage of water can be calculated as:

    percent water is equal to mass of water over mass of hydrate times 100 
    SAFETY PRECAUTIONS
    1. Always wear chemical splash goggles when working in this lab.
    2. Gloves are provided, if you choose to use them.
    3. Be sure all notebooks, book bags, purses, etc., are not placed near open flames.
    4. Be careful when handling hot glassware.
    5. Extinguish flames when not in use.
    6. Clean your work area, and place all waste/equipment in the appropriate place when finished.
    7. Wash your hands once you leave the lab.

    EQUIPMENT* AND CHEMICALS NEEDED

    Table \(\PageIndex{1}\): Equipment and Chemicals
    Equipment Equipment and Chemicals
    Milligram balance Watch glass
    spatula 250 mL beaker
    Ring stand Matches, lighter, or Striker
    Ring clamp Beaker tongs
    Wire gauze/gauze pad Bunsen Burner and tubing
    - Unknown- Alum (Potassium aluminum sulfate hydrate)

    * Images of equipment needed in this lab are in the appendix (the equipment may differ slightly or be subject to changes; follow your instructors’ directions).

    EXPERIMENTAL PROCEDURE

    1) Obtain the combined mass of an empty 250.0 mL beaker and a watch glass.

    2) Add approximately 1.000 g – 1.500 g of hydrate (formula of the anhydrous compound is KAl(SO4)2) to the beaker, and get the combined mass of the beaker, watch glass, and hydrate.

    3) Set up a ring stand with a ring clamp and a wire gauze/gauze pad. Place a Bunsen burner a few inches below the gauze pad. Once your setup is complete, have your instructor check it before proceeding to the next step.

    Beaker covered with watch glass containing Alum hydrate is being heated using a Bunsen burner.jpg

    4) Place the beaker containing the hydrate on the gauze pad and cover it with the watch glass. Turn on the burner (following all safety procedures presented by your instructor) and heat the hydrate. You will notice moisture forming on the beaker and the watch glass. Heat the hydrate for an additional two minutes, until no moisture remains.

    5) Extinguish the burner, and allow the beaker to cool to room temperature. It may take up to 20 minutes to cool to room temperature..

    6) Once the beaker is cool, get the combined mass of the beaker, watch glass, and anhydrous compound.

    7) Dispose of the remaining solid in the solid inorganic waste container, wash and dry your glassware, and complete steps 1-7 for a second trial.

    PRE-LAB QUESTIONS                                

    Name:____________________________________

    1) In your own words, define the following:

    Empirical formula:

     

    Hydrate:

     

    Anhydrous compound:

     

    2) 1.000 g of a magnesium sulfate (MgSO4) hydrate was heated, forming 0.489 g of anhydrous compound. Using this information, determine the empirical formula of the hydrate and the percentage of water present in the hydrate.

     

     

    DATA AND OBSERVATIONS

    Name: _________________________Lab Partner(s): ______________________________

    Table \(\PageIndex{2}\): Data Table

    Data

    Trial 1

    Trial 2

    Empirical formula of hydrate

    		    

    Mass of anhydrous compound (MAC)

    (MAC = M3 - M1)

    		    

    Mass of beaker and watch glass (M1)

    		    

    Mass of beaker, watch glass, and anhydrous compound (M3) (After heating )

    		    

    Mass of beaker, watch glass, and hydrate (M2)

    		    

    Mass of hydrate (M)

    (M = M2 - M1)

    		    

    Mass of water driven off (MH2O)

    (MH2O= M - MAC)

    		    

    Moles of anhydrous compound, KAl(SO4)2

    		    

    Moles of water

    		    

    Show all of your calculations here for both trials:

     
     

    POST-LAB QUESTIONS

    1. List some potential sources of error in determining the empirical formula of the hydrate.

       
    2. What is the average percentage of water in the KAl(SO4)2 hydrate? (Show calculation.)

       
    3. A hydrate of Na3B4O7 contained 62.23% water by mass. Using this information, determine the empirical formula of the hydrate. (Hint: Assume 100 g of compound.)

       

    Please click here to access the Pre-Lab, Data Tables, and Post-Lab in Word or PDF format. Complete them and upload according to your instructor's instructions.  

    This page titled 8: LAB 8 - EMPIRICAL FORMULAS AND HYDRATES is shared under a not declared license and was authored, remixed, and/or curated by Saadia Khan.