Cliffs AP Chemistry, 3rd Edition

Experiment 2: Determination of the Percentage of Water in a Hydrate

Background:Many solid chemical compounds will absorb some water from the air over time.
In many cases, this amount is small and is only found on the surface of the crystals. Other
chemical compounds, however, absorb large amounts of water from the air and chemically bind
the water into the crystal structure. The majority of these compounds are ionic salts (metallic
cation other than H+and nonmetallic anion other than OH–or O2–). To remove the water from
many of these hydrates, one only needs to gently heat the compound to slightly above the tem-
perature of boiling water. When one heats the hydrate, the crystalline structure will change and

often a color change occurs. For example, CuCl 2 ⋅2 H 2 O, copper (II) chloride dihydrate is

green as the hydrate; as the anhydride, CuCl 2 , it is brownish-yellow. When hydrates lose water
spontaneously, they are said to effloresce. The degree of efflorescence is a function of the rela-
tive humidity. Some compounds will spontaneously absorb water from the air. These com-
pounds are known as desiccants and are said to be hygroscopic (hydrophilic). When desiccants
absorb so much water that they dissolve, they are said to be deliquescent. Do not confuse pro-
duction of water vapor as an absolute indicator of dehydration since many organic compounds
such as carbohydrates will produce water vapor upon decomposition. This decomposition is not
reversible but is usually reversible in the case of an anhydride (that is, adding water to the an-
hydride will convert it back to the hydrate). In general, the proportion of water in the hydrate is
a mole ratio (whole integer or a multiple of^1 ⁄ 2 ) to that of the salt.

Scenario:A student placed some blue CoCl 2 on a watch glass and observed that over time it
changed to violet and then to red. Upon heating the red compound gently, she observed that
it changed back to violet and then to blue. She then took a sample of sodium sulfate decahy-
drate whose large crystals appeared colorless and transparent and observed over time that they
changed to a fine white powder. She then added water to the fine white powder, dissolving it.
Upon gently heating this mixture, the large, colorless transparent crystals reappeared. Then, she
took 1.700 g of a green nickel sulfate hydrate and heated the hydrate completely in a crucible
(see Figure 1).

Figure 1

Heat with cover closed to

facilitate reactions and for

reactions requiring minimal

amounts of oxygen.

Triangle

Pipe Stem

Ring,

Support

Stand,

Support

Laboratory

Burner

To prevent splattering of

contents, heat gently.

Heat with the cover open for

drying precipitates, ashing, etc.

Part III: AP Chemistry Laboratory Experiments