Hydrate Lab

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  • Topic: Magnesium sulfate, Water, Hydrate
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  • Published : April 7, 2013
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The % Composition and Empirical Formula of a Hydrate: It Doesn’t Hold Water, Or Does It?

1. To determine the percent water in an unknown hydrate. 2. To calculate water(s) of crystallization for an unknown hydrate. 3. To determine the formula of an unknown hydrate.



Proper use of the following equipment: Dial-O-Gram balance (Laboratory Technique I), electronic balance (Laboratory Technique II) and Bunsen burner (Laboratory Technique III).

Dial-O-Gram balance, electronic balance, Bunsen burner and hose, striker, ring stand, small iron ring, clay triangle, crucible and cover


Various hydrates

Hazard Alert
Avoid breathing dust or skin contact with hydrate. Clean up all spills immediately with a wet sponge or foam brush.

If an opinion is said to not hold water it means that the point of view or statement put forward is illogical, inadequate, not sound or can be shown to be wrong. Hydrates are inorganic salts that hold water. They contain a specific number of strongly bonded water molecules as part of their crystal structure. Such water molecules are called water(s) of hydration, or water(s) of crystallization. The formula of a hydrate consists of the formula of the anhydrous (without water) compound followed by a dot, then the specific number of water molecules. The dot in the formula indicates a type of bond, while strong, can usually be broken with the application of moderately high temperature. An example of such a formula is CaSO4•2H2O, commonly known as gypsum. Here, the anhydrous compound is CaSO4 and 2 in the formula represents the specific number of water molecules for each unit of CaSO4. Some anhydrous salts will pick up water readily on exposure to moisture in their environment. These salts are termed hygroscopic. Some of these hygroscopic substances are able to absorb so much moisture that they actually dissolve themselves and form a solution. These salts are called deliquescent. Some examples of salts containing water(s) of hydration and their decomposition temperatures are shown below. Notice how their names contain Greek prefixes representing the specific number of water molecules in the hydrated compound. Decomposition Temps. -6H2O (150 °C) -7H2O (200 °C) -8H2O (60 °C) -10H2O (320 °C) -5H2O (100 °C)

Figure 1 A leaky pail does not hold water

Formula MgSO4•7H2O

Name Magnesium sulfate heptahydrate (Epsom salts) Sodium tetraborate decahydrate (borax) Sodium thiosulfate pentahydrate

Na2B4O7•10H2O Na2S2O3•5H2O

Table 1: Hydrates and their decompostion temperatures.

As can be seen from Table 1, water(s) of hydration can be driven off by heating. In some cases, as with sodium tetraborate decahydrate, gentle heating will drive off only some of the water molecules while more vigorous heating is necessary to drive them all off. An example of a chemical equation showing H2O being driven off by heating appears below: MgSO4•7H2O (hydrated salt) heat MgSO4 + 7H2O

(anhydrous salt) + (water vapor)

Some hydrates spontaneously lose water molecules to the environment without the application of heat. This process is called efflorescence and these hydrates are called efflorescent. We can calculate the percent water in a hydrate by determining the amount of water driven off when a known mass of hydrate is heated. First we weigh an empty cruible and cover. Then we put a sample of the hydrate into the crucible, cap it with the cover and reweigh. The difference in weight will give us the mass of the hydrate. We heat the crucible to drive off the water. To make sure that all the water is driven off, we do a series of heating, cooling, and weighing until two consecutive weighings are the same (± 0.005 g). Such a process is called heating to constant mass. When all the water is driven off, we weigh the crucible, cover and anhydrous salt. The difference between the crucible and cover and the crucible, cover and anhydrous salt gives us the...
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