Many chemical compounds, even in relatively dry form (like powder or crystals) contain water molecules in a definite proportion to the main molecule.
For example, cobalt chloride, CoCl2, can come in two forms. Anhydrous CoCl2 contains no water (an-hydrous literally means "without water"), and is a light blue powder.
But you are more likely to find cobalt chloride in its hydrated form, CoCl2·6H2O, which means that for every six CoCl2 units, six – and exactly six – water molecules are also present. CoCl2·6H2O is a deep fucshia-colored, translucent crystal. When anhydrous CoCl2 is left out in humid air, it will rapidly soak up water and form the fucshia crystals. Both forms are pictured below.
Hydrates are substances that tend to associate with a fixed number of intact water molecules. Hydrates are almost always the (or one of the) cystal form(s) of the substance.
We indicate a hydrate by appending the dot ( · ) symbol followed by the number of water molecules, like AxBy· nH2O, where n is the number of waters.
We use the prefixes mono, di, tri, tetra, penta, hexa, hepta, octa, nona and deca when we say the names of thes compounds, like
Like everything in chemistry, the reason that many ionic compounds tend to associate with a specific number of water molecules is energy. Nature always seeks to minimize energy.
Many salts are more stable when they organize into regular crystalline structures, structures where each ionic compound is situated in a regular, prectable way. Many sugch crystals can't be formed unless a few water molecules form bridges between molecules by hydrogen bonding (see water).
The figure on the right shows a part of the crystal structure of CuSO4·5H2O, copper (II) sulfate pentahydrate. The top figure shows a part of the repeating unit of the structure; only three of the five waters are shown for clarity. Notice that in the crystal, the waters form bridges between the regularly-positioned CuSO4 compounds.
If you remember a little about energetics, you will recall that any spontaneous event, like spontaneous formation of a hydrated crystal, must have a negative Gibbs energy change, ΔG. Recall that ΔG = ΔH - TΔS, where ΔH is the enthalpy change and ΔS is the entropy change.
Because the crystal has much less entropy (disorder) than an amorphous (unstructured or chaotic) solid, ΔS is negative. That means that crystallization of CuSO4·5H2O must be "enthalpy driven," and it is. ΔH is negative and large enough to make up for the loss of entropy, thus the process is spontaneous.
Hydrates are usually more thermodynamically stable than anhydrous salts, so they form spontaneously when anhydrous salts are exposed to water, like atmospheric humidity. Waters of hydration help hydrated salts form well-ordered crystals.
Anyhydrous ionic compounds (salts) tend to be hygroscopic; they very readily soak up water from the surroundings, like from atmospheric humidity.
When anhydrous salts are opened in a laboratory, they immediately begin to hydrate with atmospheric water. Often anhydrous salts are kept in dry environments that exclude water, like under a dry argon (Ar) atmosphere (it turns out to be very easy to extract water from argon gas, plus it's abundant and cheap).
When using hydrated salts, we have to be careful to incorporate the waters of hydration when calculating the formula weight. For example, the formula weight of copper (II) sulfate is 159.6 g/mol, but the FW of CuSO4· 5H2O is 249.7 g/mol.
When you work with anhydrous salts in an open lab, the formula weight of the compound will change over time as it absorbs water.
Finally, you might think that an anhydrous salt would have the same solubility as its hydrated counterpart. But a given anhydrous salt can have very large lattice energy, the energy that holds the molecules together in the solid. Thus, they might not ionize readily in water. Some anhydrous salts will just clump together and float on top of water, dissolving only very slowly.
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