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Formula of a Hydrate

Parts of the text on this lab will not print out. This is by design. The parts that won’t print are notes for teachers. Students don’t need those notes and they are automatically excised from the printout.

Objective

Note for Teachers: Give out this lab handout at least one day before doing this lab and assign the problems for homework.

The lab work has two objectives: first, confirm the formula of a hydrate with known formula and second, find the formula of a hydrate in which the salt formula is known but not the molar amount of water.


Overview

This lab activity combines in-class problem-solving with lab work. Students will be introduced to the concept of hydrated salts, anhydrous salts, hydrate nomenclature, and the molar mass of hydrates. The lab portion of this work conists in confirming the formula of a hydrate of known formula (CuSO4·5H2O) and finding the formula of a hydrated salt of known formula but unknown water content.


Background

Some chemical compounds, especially inorganic salts, incorporate water into their crystalline structures. Water has a polar structure: it has positively and negatively charged parts within each molecule. This gives it a strong attraction toward ions. The ions in some salts attract and form strong bonds with water molecules. These salts, when they have absorbed water, are called hydrates. Anhydrous salts are salts that can form hydrates but which have had all the water driven off, usually by heat. Hydrated salts are characterized by the number of moles of water molecules per mole of salt. The so-called water of hydration of nickel (II) chloride (NiCl2) is six moles H2O for every one mole of NiCl2. The hydration reaction is shown below. The hydrate in this reaction is called nickel (II) chloride hexahydrate.

NiCl2 + 6H2O Arrowsngl NiCl2·6H2O

The formula of this hydrate shows the molar amount of water incorporated into the crystal matrix. For most hydrates the amount of water included in the formula is only important when trying to measure molar amounts of the salt. You need to know the true formula weight (molar mass) in order to measure out the mass needed to give a certain number of moles. The chemical importance of the water of hydration is minimal since it can be driven off by heat or simply dissolve away if the salt is dissolved in water. From the reaction above, nickel (II) chloride hexahydrate (NiCl2·6H2O) the molar mass is 237.69 g/mol not 129.60 g/mol. The figure 129.60 g/mol is the molar mass of the anhydrous salt.

Greek Prefixes
  1. Mono-
  2. Di-
  3. Tri-
  4. Tetra-
  5. Penta-
  6. Hexa-
  7. Hepta-
  8. Octa-
  9. Nona-
  10. Deca-

Formulas for hydrates are written using a dot convention: a dot is used to separate the formula of the salt from the formula of the water of hydration. A numerical coefficient gives the molar amount of water included in the hydrate. Hydrates are named using prefixes for the word hydrate (at right). For example, CuCl2·2H2O is copper (II) chloride dihydrate and CuSO4·5H2O is copper (II) sulfate pentahydrate. One key point: the dot is not a multiplication sign. When calculating the molar mass you add the molar mass of water (multiplied by the coefficient).

An everyday example of hydration is concrete. Concrete is made by mixing Portland cement with water and aggregate materials. The aggregate materials are the gravel and sand that add strength to the final concrete. The Portland cement is a mixture of calcium silicates, calcium aluminate, calcium aluminoferrite and gypsum. All of these chemicals absorb water by hydration. This means that concrete does not ‘dry’ in a conventional sense. Instead the water mixed with the concrete combines chemically with the materials in the cement and the resulting hydrates form a strong matrix that holds the concrete together and makes it strong.

Another interesting example of the value of hydration is the incorporation of hydrated building materials (such as concrete, gypsum wall board and plaster). The building materials will not rise above the 100°C boiling point of water until all of the water of hydration has been driven off. This can help keep damage to a minimum until the fire can be put out. In the construction business this is known as passive fire protection.




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Pre-lab Problems

The following problems will help you to be able to do the math required for the analysis of your lab results.

Find the molar mass of the following hydrated and anhydrous salts.
  1. FeCl2·2H2O ____________
  2. FeCl2 ____________
  3. CaCl2·2H2O ____________
  4. CaCl2 ____________
  1. CuSO4·5H2O ____________
  2. CuSO4 ____________
  3. MnSO4·4H2O ____________
  4. MnSO4 ____________
Example:
Given 0.62 g CuSO4·5H2O find the mass of water that would be driven off by heating
          1 mol
0.62 g · ——————— = 2.48 × 10-3 mol          reaction: CuSO4·5H2O --> CuSO4 + 5H2O 
         249.69 g
                   5 mol H2O 
2.48 × 10-3 mol · ————————————— = 1.24 × 10-2 mol H2O = 0.22 g H2O (18.02 g/mol)
                 1 mol CuSO4·5H2O
So the anhydrous salt in the sample accounts for 0.40 g 
and the mass of the water of hydration is 0.22 g
  1. Find the mass of the anhydrous salt in a 142.3 g sample of MnSO4·4H2O. This might be found in the lab by heating the sample until its mass does not decrease any further.
  2. Find the mass of the anhydrous salt and the mass of water in a 10.9 g sample of FeCl2·2H2O.
Example:
A 140.5-g sample of NiSO4·XH2O is heated until no further decrease in mass.
The mass of the anhydrous salt is 77.5 g. Find the number of water molecules 
in the formula of this hydrate of nickel (II) sulfate.
          1 mol
77.5 g · ——————— = 0.50 mol NiSO4         reaction: NiSO4·XH2O --> NiSO4 + XH2O 
         154.76 g
                               1 mol
140.5 g - 77.5 g = 63 g H2O · ——————— = 3.5 mol H2O
                              18.02 g 
3.5 mol H2O
—————————— = 7 mol H2O per mol anhydrous salt so formula is NiSO4·7H2O
  0.5 mol
nickel (II) sulfate heptahydrate



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Problems, cont.

  1. Given that a 40.14-g sample of hydrated NiSO4 is reduced in mass to 22.14 g upon heating show that the formula of the hydrate is NiSO4·7H2O.
  2. Given that a 139.4-g sample of hydrated MnSO4 is reduced in mass to 94.38 g upon heating find the empirical formula of the hydrate. Also, write the name of the hydrate of manganese sulfate.

Materials

  1. ring stand
  2. ring
  3. crucible tongs
  4. crucible
  5. clay triangle
  6. small dropper for water

  1. bunsen burner
  2. matches
  3. lab balance
  4. lg watch glass
  5. ~1 g CuSO4·5H2O
  6. 2 - 3 g of an unknown hydrate

Safety

Procedure

Part I

  1. Thoroughly heat your crucible and allow it to cool on the clay triangle. Then weigh your crucible. This ensures that it is clean and dry.
  2. Obtain about 1 g of CuSO4·5H2O in your crucible and weigh it, recording the mass to the maximum available precision.
  3. Make a prediction about the amount of mass that will be lost when you heat the sample of copper (II) sulfate pentahydrate. Specifically, calculate the mass of the anhydrous salt and the mass of water that will be driven off. Show this to your instructor before proceeding.
  4. Once you have your instructor’s approval, place the crucible containing the CuSO4·5H2O on the clay triangle.
  5. Light the bunsen burner and adjust for a hot flame.
  6. Heat the crucible as gently as possible with the burner by moving the burner under the crucible for a few seconds at a time. Note the release of any steam from the crucible.
  7. Remove the heat source and use a pair of lab spatulas to occasionally stir the copper sulfate. Carefully scrape all of it back into the crucible. Be careful not to do this while heating!
  8. Continue heating gently until the salt turns completely white. Be careful not to overheat! The heat can become so intense that the sulfate in the salt begins to break down. If this happens the salt will turn yellow and produce a sulfurous smell. It will also ruin your data since it will reduce the mass more than expected due to the decomposition of the salt.
  9. Stop heating when the salt has lost all traces of blue color. Allow the crucible and its contents to cool completely.
  10. Once the crucible is cool, find its mass. Then stir the copper(II) sulfate and heat the crucible and its contents again for a short time. Then allow it to cool and weigh it. If the mass is the same as the previous weighing, then the salt has been completely dehydrated. If not, repeat the heating/cooling/weighing process until two successive weighings have the same mass.
  11. Record the final mass of the anhydrous salt in you lab notebook and do the calculations to show that the molar ratio of water to anhydrous salt really is 5:1.
  12. Empty the anhydrous salt onto a large watch glass. Use the dropper to add a very little water to the anhydrous copper (II) sulfate. Describe what happens in your lab notebook. For your report think about what is happening at the molecular level when you add water. Draw a model.
  13. When you finished this part of the lab empty the re-hydrated CuSO4·5H2O into the beaker provided by your instructor for this purpose. Then begin Part II.

Part II

In this part of the lab you will repeat the same procedure performed for the salt of known formula with a salt for which you do not know the hydrate formula. The salt may be magnesium sulfate (MgSO4), sodium phosphate (Na3PO4), calcium chloride (CaCl2), or sodium carbonate (Na2CO3). To ensure better chances of getting the correct result you may want to consider doing at least two (and perhaps three) trials. For each trial use a minimum of 2 g but no more than 3 g. These salts are not nearly so hazardous as the copper (II) sulfate and a larger amount will help to reduce errors due to small lab balance inaccuracies. For all of these salts, both the hydrate and anhydrous salt are white. Some of them release so much water that if you let it boil it will spatter all over you and the table. Be careful! Finally, unless you frequently stop heating to stir the crystals they will combine and harden, possibly trapping water inside. One method which may possibly prevent this is to grind the hydrated salt in a mortar before heating it. Weigh the hydrate after you grind it if you do grind it up.


The Report

For this lab simply provide neat data tables, sample calculations, and the answers to the following questions in a professional-quality typed document.

  1. In a data table give the following information: starting mass of hydrated copper(II) sulfate, expected mass of anhydrous copper (II) sulfate, and expected mass of water to be lost by heating. Also, show your calculations for the expected masses.
  2. In a data table give the following information: starting mass of hydrated copper(II) sulfate, final mass of heated (anhydrous) copper(II) sulfate, mass of water lost. How well did your prediction match up with your results for copper (II) sulfate pentahydrate?
  3. Say that your mass of water lost was too low compared with your prediction. What may have caused this? Describe a scenario as it could really have happened in the lab.
  4. Say that your mass of water lost was too large compared with your prediction. What may have caused this? Describe a scenario as it could really have happened in the lab.
  5. Report data for your unkonwn hydrate in the same fashion as you did for copper(II) sulfate in the second question above. Show work for your calculations. What is the formula of your hydrate? What is the name of this hydrate?
  6. Look up the correct hydrate formula online. Did you find the correct formula? If so, explain how you avoided errors. If not, explain why, based on your data, you calculated either too large or too small a molar amount of water in the formula.
Last updated: Sep 05, 2018 Home