Determining a Mole Ratio: Hydrate Analysis Lab
You will use techniques of quantitative analysis to determine the percent by mass of water in a solid and the formula for water of hydration (or hydrate). You will actually be determining the mole ratios of the products through mass analysis. The goals of this experiment are:
(1) to perform the student’s first stoichiometric calculation using the data collected in the laboratory
(2) to improve the accuracy using balances which have greater sensitivity
(3) to obtain the skill of manipulation of crucible and tongs
(4) to observe a crucible coming to constant weight (mass)
(5) to demonstrate the law of conservation of mass
(6) to recognize that ionic crystalline solids have water molecules trapped in their geometric structure
(7) to calculate the formula of a hydrate salt unknown if the anhydrous formula is known.
The accurate determination of the composition of substances is essential in modern technology and important in our everyday life. Typical problems in which analytical chemistry plays a role are determining the compositions of lunar and other extraterrestrial samples, monitoring harmful levels of pollutants in the atmosphere above our cities and in our water supplies, maintaining quality control over drug and food products, detecting trace impurities in ultra pure semi-conductor materials (transistors, diodes, etc.), and making clinical investigations which determine the nature and concentration of materials in biological fluids.
As can be seen from the examples given, analytical determinations may be either qualitative or quantitative in nature. A qualitative analysis is aimed at determining the identity of a substance (what is present). The previous cation and anion labs were qualitative in nature. A quantitative analysis determines the amounts of known substances present in a particular sample (how much is present).
This experiment will promote you from using the top loading balances on the lab benches (0.01g) to improved precision weighing and quantitative analysis techniques. You will use the one three decimal (.001g) top loader and/or the analytical balance (0.0001g).You will determine the mole ratio of a hydrate decomposition reaction by removing a mass of water in a sample of solid hydrate.
In the 5th edition of the Corwin text Section 13.10 introduces the concept of
a hydrate . In this lab we
will do a simple calculation to determine the Percent Water in a known and an unknown hydrate. This calculation
will be at least three significant, if not four significant figures. Then using
the same data we will calculate the water of hydration through determining the
mole ratio of the decomposition reaction. The data will be translated into a
one significant figure answer.
A hydrate is a crystalline compound that contains a specific number of water molecules attached to an ionic formula unit.
Chemists know that particles of a crystalline solid are
arranged in a regular geometric pattern. However, the particles can be of
different types. They can be ionic,
molecular, or metallic. In Ionic Solids the crystals are composed of
regular patterns of ions.
Refer to Corwin’s section 13.5 for more discussion
of Molecular and Metallic Solids.
Water molecules are part of the crystalline structure in hydrates. The formulas for hydrates are written symbolically: MgSO4∙7H2O which is Epsom salts. The ∙ is read “in combination with”. This formula is read “Magnesium sulfate in combination with seven waters of hydration”.
When there is no water of hydration, the term Anhydrous is used. In some cases, the hydrate can form spontaneously from the anhydrous salt if sufficient moisture is present in the air.
For example, Anhydrous Calcium chloride readily absorbs water molecules from the air to form its common hydrate: CaCl2∙2H2O and is used in lab desiccators for drying agents. Sometimes gases are routed through a tube of Anhydrous Calcium Chloride to remove the water molecules from the gas. The reaction is:
CaCl2 (s) + 2 H2O (g) → CaCl2∙2H2O (s)
Here is a reaction, maybe you did when you were a kid (or something similar if you broke a limb and had to have a cast). (CaSO4)2∙H2O is ‘Plaster of Paris’.
Did you ever make a hand print
when you were a kid?
Plaster is widely used as a support for broken bones; a bandage impregnated with plaster is moistened and then wrapped around the damaged limb, setting into a close-fitting yet easily removed tube, known as an orthopedic cast; however, this is slowly being replaced by a fiberglass variety.
Sometimes ‘Plaster’ is written with a fraction: CaSO4∙½H2O. There are two calcium ions, two sulfate ions for every molecule of water. (Many chemists do not like using the fractions in formulas-but both are basically the same formula).
Plaster of Paris is a white powder. You mix it with water and it sets as a hardened substance after it absorbs water and changes its crystalline structure.
(CaSO4)2∙H2O (s) + 3 H2O (l) → 2 CaSO4∙2H2O (s)
The hardened product CaSO4∙2H2O is known as gypsum. Sheetrock has gypsum under the outside paper coating. If you heat gypsum at 150ºC the reverse reaction will occur.
2 CaSO4∙2H2O (s) → (CaSO4)2∙H2O (s) + 3 H2O (g)
The water comes out of the gypsum as steam. We are going to do this type of reaction in this experiment. We will carefully heat the hydrate in a crucible and drive off the molecules of water to form either another hydrate with less water or totally push the crystal to its anhydrous form. The geometric structure of an anhydrate is different from the hydrate of the crystal.
Our objective is to determine the hydrate formulas of one of the following:
1 CuSO4∙XH2O (s) → 1 CuSO4 (s) + X H2O (g)
1 CoCl2∙XH2O (s) → 1 CoCl2 (s) + X H2O (g)
1 NiSO4∙XH2O (s) → 1 NiSO4 (s) + X H2O (g)
1 BaCl2∙XH2O (s) → 1 BaCl2 (s) + X H2O (g)
It is NOT usually obvious from an object’s appearance that it is HOT! Severe burns can result if you are not very careful when handling a Burner, tripod, clay triangle, or crucible when these are HOT!
Many hydrates dissolve in their water of hydration on initial heating forming a concentrated solution which easily spatters and sometime violently (like popping pop corn). Therefore, heat gently initially and then strongly only after most of the water has been driven off. In the first three unknowns above there is a significant color change so you will know the extent of your reaction after initial heating.
CAUTION: The solids may be poisonous. Wash skin contact areas profusely.
Equipment: Hardware: Instruments:
(from cart) (from lab stations) (instructor’s desk)
Crucible & Cover Burner & Tubing Top Loading Balance (0.001g)
Clay Triangle Analytical Balance (0.0001g)
Tripod (or Ring Stand and 5” ring)
Solid Chemical Hydrates (all not required):
Barium Chloride Copper II Sulfate
Calcium Sulfate Nickel II Chloride
Cobalt II Chloride Nickel II Sulfate
Copy the data page into your lab notebook.
Part I: Bring the crucible to Constant Weight:
1. Clean as best you can a crucible and cover. (If the crucible has a black or colored stain, heat gently with a little dilute nitric acid to dissolve residues which are soluble-we will skip this step as we are out of nitric acid Spring 2009)
2. Weigh the crucible and cover on the Top Loader at your desk. Record the mass to the nearest 0.01 g. Then weigh it on the 0.001g top loader in the front of the room. Record these masses on your data page in your lab notebook. After this weighing, you should never touch the crucible until the end of the experiment when you have recorded the final weighing.
3. Place a clay triangle on the tripod (or on the 5” ring attached to the ring stand.). (If using the ring stand adjust the ring so that the bottom of the crucible will be one or two inches above the burner.)
4. Practice with your crucible tongs inserting the crucible (without lid) into and out of the clay triangle. Then try it with the lid on after you have achieved the needed skills.
5. Heat the crucible and cover using the Fisher burner until the bottom becomes dull red. Ajar the lid to allow moisture inside to escape as shown in Figure A. Heat for an additional minute after the crucible becomes red.
6. Discontinue heating and let the crucible cool to room temperature. Wait between 7 and 9 minutes for the cooling. The cooling time between each weighing should be within a minute or two so that if the crucible is still warm, your masses will not vary because of the temperature. Note: a warm crucible will heat the air inside the crucible and immediately around the crucible so the expanded air weighs less due to Charles Law of gases.
7. Weigh the crucible on the 0.001g Top Loader, then on the Analytical Balance (0.0001g) in the front of the room. Record the masses on your data page. If the masses are within 2 mg (0.002 g) of the initial preheating weighing, then you may proceed to Part II. If not, heat the crucible and cover one more time and check the mass. When within 2 mg, you have achieved constant mass of the crucible and cover.
Part II: Analysis of the Hydrate Unknown
8. Obtain from your instructor a Hydrate Unknown. Your instructor will tell you the formula for the Anhydrate.
9. Using your top loading balance on your desk, place between one and two grams of the hydrate into your crucible. Then record the mass using either the 0.001g Top Loader on the front desk or the 0.0001g Analytical Balance. For the remainder of the experiment always use the same balance either the 0.001g or the 0.0001g.
10. Using the Bunsen Burner heat the crucible gently. Have the lid slightly open (or ajar) as shown in Figure A.
11. Increase your heating. If white smoke or a cracking noise is heard, slow down the heating. Gradually increase the heating until the bottom of the crucible is a dull red. Shut off the burner and allow the crucible to cool.
12. After sufficient cooling, reweigh. Check the contents for a color change if it is one of the highly colored hydrates that have a different color as an anhydrate. Record the mass on your data page.
13. Heat the crucible with the Fisher burner until the bottom of the crucible is dull red. Shut off the burner. Allow the crucible to cool.
14. Weigh the crucible and contents. If the crucible and contents are within 2 mg (0.002g) of the previous weighing you are finished. Otherwise repeat the heating and cooling cycle until the crucible and content achieve constant weight.
15. When finished add 1 drop of DI water to the contents. Note the instant color change if it is one of the highly colored hydrates. Did the color return to the original color of the crystals before any heating?
16. You may clean the crucible and dispose of the contents as described by your instructor. There should be special waste containers for each Unknown. These crystals may be dried and used again in the future.
Part III: Calculation of the Hydrate Formula
17. Calculate the molar mass of the anhydrous salt using the periodic table.
18. Determine the grams of anhydrate remaining in the crucible.
(Subtract the Constant Weight of the Empty Crucible from the Constant Weight After Heating.)
19. Calculate the number of moles of anhydrate.
(Divided the grams of anhydrate by its molar mass to find the number of moles.)
20. Determine the number of grams of water which escaped from the crucible. (Substract the Constant Weight of the Crucible and Anhydrate from the Mass of the Crucible and contents before heating.)
21. Determine the molar mass of water from the Chart.
22. Calculate the number of moles of water which evolved.
(Divide the mass of water by its molar mass.)
23. Determine the mole ratio of the water to the anhydrous salt.
(Divide the moles of water by the moles of anhydrous salt.)
24. Determine the formula of the hydrate. Round off the mole ratio to the nearest whole number if less than 0.19 or greater than 0.81 fractions. Otherwise, multiply the mole ratio by a number which gives two whole numbers for the ratio (within 0.1 range of the whole number). A 2.5 to 1 is really a 5 to 2 ratio.
25. Also determine the % water in the hydrate.
(Divide the mass of the water by the mass of the hydrate and multiple by 100 to create the % unit).
25. Have your instructor initial you data page before you leave the lab.