Many solid chemical compounds incorporate water molecules directly into their internal architecture. These substances, known as hydrates, are crystalline solids where the water molecules are present in a fixed ratio relative to the compound itself. This integral water is often called the water of hydration, or sometimes the water of crystallization. This water is not merely stuck to the surface, but is a fundamental part of the compound’s chemical composition.
Defining Water of Hydration and Hydrates
The water of hydration represents water molecules that are chemically integrated into the crystal lattice structure of an ionic compound. Unlike absorbed moisture, this water is bound to the central ions or molecules through coordination bonds or strong intermolecular forces, ensuring it maintains a precise, stoichiometric ratio to the host compound. This fixed proportion means the water molecules are essential for forming the specific geometric shape and physical properties of the hydrate.
Consider the well-known example of copper(II) sulfate pentahydrate, with the formula \(\text{CuSO}_4 \cdot 5\text{H}_2\text{O}\). In this deep blue crystal, four of the five water molecules are directly coordinated to the central copper(II) ion, forming a complex ion within the solid structure. The fifth water molecule is present in the lattice, held by hydrogen bonds to the sulfate ion and the other water molecules. This structural arrangement confirms that the water molecules are necessary to stabilize the entire crystalline network.
The compound without its water component is referred to as the anhydrous form, and its properties are distinctly different from the hydrated form. For instance, a hydrate like magnesium sulfate heptahydrate (\(\text{MgSO}_4 \cdot 7\text{H}_2\text{O}\)), commonly known as Epsom salt, exhibits a precise crystalline form. The presence of the water molecules dictates the spacing and arrangement of the ions, meaning that the hydrate and its anhydrous counterpart are technically different chemical species with unique structures.
Naming Conventions for Hydrates
The chemical nomenclature for hydrates indicates the exact number of water molecules bound within the crystal structure. The naming process starts with the name of the ionic compound (the anhydrous portion). This is followed by a term specifying the number of water molecules, ending with the word “hydrate.”
A Greek prefix is attached to the word “hydrate” to denote the precise mole ratio of water to the anhydrous salt. For example, a compound with one water molecule is a monohydrate, two is a dihydrate, three is a trihydrate, and five is a pentahydrate. The chemical formula for a hydrate uses a centered dot to separate the anhydrous compound from the water molecules, such as \(\text{CoCl}_2 \cdot 6\text{H}_2\text{O}\).
This formula represents cobalt(II) chloride hexahydrate, indicating six water molecules for every one formula unit of cobalt(II) chloride. Other examples include calcium sulfate dihydrate (\(\text{CaSO}_4 \cdot 2\text{H}_2\text{O}\)), known as gypsum, and iron(II) chloride tetrahydrate (\(\text{FeCl}_2 \cdot 4\text{H}_2\text{O}\)). This systematic nomenclature ensures the chemical composition, including the water content, is unambiguous.
The Dehydration Process
The water of hydration is removed through dehydration, typically by heating the hydrate. Thermal energy overcomes the forces holding the water molecules within the crystal lattice, causing the water to escape as vapor. The resulting solid substance left behind is the anhydrous compound, meaning “without water.”
This removal of water often results in a dramatic change in the compound’s physical characteristics, such as color and crystal shape. For example, the bright blue crystals of copper(II) sulfate pentahydrate transform into a white, powdery residue of anhydrous copper(II) sulfate upon heating. The loss of the integral water molecules causes the crystalline structure to collapse, leading to a new substance with different properties.
While heating is the deliberate method for dehydration, some hydrates spontaneously lose water to the atmosphere in a process called efflorescence. This happens when the vapor pressure of the water in the hydrate exceeds the partial pressure of water vapor in the surrounding air. Conversely, some anhydrous compounds are hygroscopic, readily absorbing moisture; if they absorb enough water to dissolve, the process is called deliquescence.
Practical Applications of Hydrates
The ability of compounds to incorporate and release water molecules makes them valuable in industrial and laboratory settings. A primary application is their use as desiccants, or drying agents. Anhydrous salts, such as calcium sulfate or magnesium sulfate, are effective desiccants because they are hygroscopic and absorb moisture to form stable hydrated structures.
This property is utilized to remove trace amounts of water from solvents, gases, or sealed environments like desiccators, which require low humidity. Many desiccants are regenerative, meaning they can be heated to remove the absorbed water and then reused. The distinct color change that accompanies hydration and dehydration also allows certain compounds, such as cobalt(II) chloride, to serve as humidity indicators.
Anhydrous cobalt(II) chloride is blue, but when it absorbs moisture to form the hexahydrate, it turns pink, providing a visible signal of humidity levels. Furthermore, forming stable hydrates is a practical way to store and handle certain chemicals, as the hydrated form is often less reactive or safer than the anhydrous form. For example, the widespread use of gypsum in construction relies on the stability of calcium sulfate dihydrate.