Water of crystallization

In molecular formulas water of crystallization is indicated in various ways, but is often vague. The terms hydrated compound and hydrate are generally vaguely defined.

Some hydrogen-bonding contacts in FeSO₄^.7H₂O. This metal aquo complex crystallizes with water of hydration, which interacts with the sulfate and with the [Fe(H₂O)₆]²⁺ centers.

A salt with associated water of crystallization is known as a hydrate. The structure of hydrates can be quite elaborate, because of the existence of hydrogen bonds that define polymeric structures.[3] [4] Historically, the structures of many hydrates were unknown, and the dot in the formula of a hydrate was employed to specify the composition without indicating how the water is bound. Examples:

• CuSO₄ • 5H₂O - copper(II) sulfate pentahydrate
• CoCl₂ • 6H₂O - cobalt(II) chloride hexahydrate
• SnCl₂ • 2H₂O - tin(II) (or stannous) chloride dihydrate

For many salts, the exact bonding of the water is unimportant because the water molecules are labilized upon dissolution. For example, an aqueous solution prepared from CuSO₄ • 5H₂O and anhydrous CuSO₄ behave identically. Therefore, knowledge of the degree of hydration is important only for determining the equivalent weight: one mole of CuSO₄ • 5H₂O weighs more than one mole of CuSO₄. In some cases, the degree of hydration can be critical to the resulting chemical properties. For example, anhydrous RhCl₃ is not soluble in water and is relatively useless in organometallic chemistry whereas RhCl₃ • 3H₂O is versatile. Similarly, hydrated AlCl₃ is a poor Lewis acid and thus inactive as a catalyst for Friedel-Crafts reactions. Samples of AlCl₃ must therefore be protected from atmospheric moisture to preclude the formation of hydrates.

Structure of the polymeric [Ca(H₂O)₆]²⁺ center in crystalline calcium chloride hexahydrate. Three water ligands are terminal, three bridge. Two aspects of metal aquo complexes are illustrated: the high coordination number typical for Ca²⁺ and the role of water as a bridging ligand.

Crystals of hydrated copper(II) sulfate consist of [Cu(H₂O)₄]²⁺ centers linked to SO₄²⁻ ions. Copper is surrounded by six oxygen atoms, provided by two different sulfate groups and four molecules of water. A fifth water resides elsewhere in the framework but does not bind directly to copper.[5] The cobalt chloride mentioned above occurs as [Co(H₂O)₆]²⁺ and Cl⁻. In tin chloride, each Sn(II) center is pyramidal (mean O/Cl-Sn-O/Cl angle is 83°) being bound to two chloride ions and one water. The second water in the formula unit is hydrogen-bonded to the chloride and to the coordinated water molecule. Water of crystallization is stabilized by electrostatic attractions, consequently hydrates are common for salts that contain +2 and +3 cations as well as −2 anions. In some cases, the majority of the weight of a compound arises from water. Glauber's salt, Na₂SO₄(H₂O)₁₀, is a white crystalline solid with greater than 50% water by weight.

Consider the case of nickel(II) chloride hexahydrate. This species has the formula NiCl₂(H₂O)₆. Crystallographic analysis reveals that the solid consists of [trans-NiCl₂(H₂O)₄] subunits that are hydrogen bonded to each other as well as two additional molecules of H₂O. Thus 1/3 of the water molecules in the crystal are not directly bonded to Ni²⁺, and these might be termed "water of crystallization".

The water content of most compounds can be determined with a knowledge of its formula. An unknown sample can be determined through thermogravimetric analysis (TGA) where the sample is heated strongly, and the accurate weight of a sample is plotted against the temperature. The amount of water driven off is then divided by the molar mass of water to obtain the number of molecules of water bound to the salt.

Water is particularly common solvent to be found in crystals because it is small and polar. But all solvents can be found in some host crystals. Water is noteworthy because it is reactive, whereas other solvents such as benzene are considered to be chemically innocuous. Occasionally more than one solvent is found in a crystal, and often the stoichiometry is variable, reflected in the crystallographic concept of "partial occupancy." It is common and conventional for a chemist to "dry" a sample with a combination of vacuum and heat "to constant weight."

For other solvents of crystallization, analysis is conveniently accomplished by dissolving the sample in a deuterated solvent and analyzing the sample for solvent signals by NMR spectroscopy. Single crystal X-ray crystallography is often able to detect the presence of these solvents of crystallization as well. Other methods may be currently available.

In the table below are indicated the number of molecules of water per metal in various salts.[6][7]

Transition metal sulfates form a variety of hydrates, each of which crystallizes in only one form. The sulfate group often binds to the metal, especially for those salts with fewer than six aquo ligands. The heptahydrates, which are often the most common salts, crystallize as monoclinic and the less common orthorhombic forms. In the heptahydrates, one water is in the lattice and the other six are coordinated to the ferrous center.[13] Many of the metal sulfates occur in nature, being the result of weathering of mineral sulfides.[14]

• 
  Hydrated copper(II) sulfate is bright blue.
• 
  Anhydrous copper(II) sulfate is white.

• Hydrate
• Mineral hydration
• Hydrous oxide