Glencoe Chapter 11
Experiment: Formula of a Hydrate

INTRODUCTION

Many salts readily dissolve in water to form solutions.  When water is allowed to evaporate from the solutions, crystals appear.  Often, the crystals that appear to be dry actually hold a good deal of water within the crystalline structure.  If the crystals are heated, however, this water is driven off.  These types of salt crystals are called hydrates.  The physical properties of a hydrate may vary from the properties of the salt without the water (called "anhydrous salts").  For instance, the color may differ.  However, the water contained in the hydrate is bound physically to the crystal of the salt and the two can be easily separated by physical means (heating).

After a hydrate is heated, the remaining salt is called anhydrous (without water).  To differentiate between the hydrate and the anhydrous salt, different chemical names and formulas are used.  

Examples: 
Hydrated copper(II) sulfate pentahydrate is written as     CuSO₄ ^. 5H₂O  
Anhydrous copper(II) sulfate is written as CuSO₄

The amount of water associated with the salt is written as a whole number ratio to the moles of salt. 
The formula CuSO₄ ^. 5H₂O states that there are 5 moles of water combined with 1 mole of CuSO₄

In this experiment, you will find the ratio of moles of water to moles of anhydrous salt for an unknown hydrate.  After determining this calculation, you will determine the formula of the hydrate.  You will then find your % error.

MATERIALS

• bunsen or fischer burner
• crucible with cover
• crucible tongs
• ring stand with screening
• striker / lighter
• centigram balance / analytical balance
• different hydrates - your teacher will assign you a specific one 
• goggles and aprons

PROCEDURE

1. Place the burner under the ring stand.  Place a clean and dry crucible with its cover onto the ring stand (with the screening to support it).  Ignite the burner using the striker.  Preheat the crucible for 60 seconds to make sure that it is completely DRY.
2. Shut off the burner and then let the crucible cool off.
3. After a few minutes of cooling, use the crucible tongs to move the crucible and then measure the TOTAL weight of the crucible and its cover.  Record this in your data table.
4. Each group will be assigned a different hydrate.  Record the letter/# assigned in the data table along with the anyhydrous molar mass for that substance (see the ChemCentral Schedule for the listing of the masses).
5. Fill your crucible so that it is 1/4 full.  Place the cover back on the crucible.
6. Measure ACCURATELY and PRECISELY the mass of your hydrate.  Record this in your data table.  (Remember, you will be calculating your ERROR at the end of this experiment so making excellent mass measurements is one of the keys to success!!!)
7. Put the crucible and the lid onto the ring stand again and GENTLY heat it for a MINIMUM of 5 minutes.  Slightly tilt one side of the lid to help ensure that the water will evaporate as you drive it off with heat.  NOTE: You don't want the compound to pop out of the crucible - if you lose the mass of the compound because it pops out, the experimental data will be useless!
8. Turn off the burner.  Cool for a couple of minutes.
9. Measure the mass of the crucible, cover and contents.  Record on data table.
10. Repeat steps 6-7 again.  Measure the mass of the crucible, cover and contents.  Record on data table.
11. Repeat steps 6-7 again.  Measure the mass of the crucible, cover and contents again.  Record on data table.
12. If the difference between the second and the third heating is 0, that means you've driven off all the water from the compound.  But, if there is more than .1 g difference, keep heating/cooling until there is no more mass change.
13. Wait until the crucible is TOTALLY cooled and then thoroughly clean out your crucible (put any crystal into the trash) and wash and dry with a paper towel.   Leave NO CONTENTS inside the crucible when you are done.  Remember, the next lab group will be using your crucible.
14. CLEANUP: Take a wet paper towel and thoroughly swab down your lab table.  Then, take a dry paper towel and thoroughly clean out the middle sink area of your lab table.  There should be "no white" streaks or areas on any black lab table area.

DATA TABLE

In the lab, be sure to use blue or black PEN (NO PENCIL) when you complete this data table.  NEVER WHITE-OUT data.  Enter data carefully so that you make no mistakes.  Print neatly!

Print out this data table!

DATA ANALYSIS QUESTIONS - These are also on the data table printout.

1. What was the final mass of the water that was driven off by heating (see calculations made on data table)?
2. Calculate the number of moles of anhydrous salt (that's the hydrate without the water).  Use these molar mass that you wrote on the top of your data table.
3. Calculate the number of moles of water that was evaporated from your hydrate (use the mass you calculated in the data table)..
4. Compare your results in question #2 to the results to question #3.  Use a ratio.

______________________moles of water
                                            moles of anhydrous salt

5. Take the moles of water and divide by the moles of anhydrous salt.  You probably will not have a whole number but round it to the nearest whole number.

Example:    0.646 moles of water / 0.106 moles anhydrous salt = 6.09 which rounds to 6.  So your ratio of water to salt is 6 : 1.  Therefore, your formula would be 1 "X" ^. 6 H₂O.

Go ahead and do the same thing for YOUR results.  Determine the whole number ratio and then write it into the correct formula using the same formatting as indicated in the example above.  Don't use "X" as in the example but instead use the correct letter for whatever hydrate YOU used.

6. Calculate your % error.  For example, let’s say you actually got 1.1 moles of water in #5.  But, remember that the moles of water will be an integer value.  So, your % error would be:

                       | 1.1 - 1 moles |    x 100   =        10% error
                           1 mole

EXPERIMENT FINAL THOUGHTS

The precision of the individual instruments used in an experiment gives a good indication of the precision of the experiment in general, assuming that the user takes advantage of all of the marked digits as well as an estimate digit for each measurement.  The accuracy of an experiment can be determined by many factors.  Carefulness in measurement, avoiding spills and splashes, using clean glassware, and following directions can all help to improve the accuracy of experimental results.  But, how is accuracy measured?  Recall that accuracy describes how close a result is to an accepted value.  One way of measuring this is by calculating percent error.  Percent error gives a numerical indication of how far the experimental results are from the actual or accepted value.  The formula used to calculate percent error is shown below:

Error =   |actual results - theoretical results|   x 100%
                           theoretical results

In some experiments, an error of twenty percent or more is considered acceptable.  In other experiments, the error can be kept to two percent or below.  the procedure and instruments can be a factor in how accurate the results are expected to be.  Of course, the lower the percent error, the greater the accuracy of the results.