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