Hydrated: The Crystalline World of Water
What is Water of Crystallisation?
Imagine building a tower with magnetic blocks. Now, imagine that you need to include specific, smaller magnetic pieces to hold the entire structure together. In a hydrated compound, the water molecules act like those smaller magnetic pieces. They are not just sitting in the gaps; they are a fundamental part of the building plan for the crystal.
The water of crystallisation is water that is chemically combined in a definite proportion within a crystal. These water molecules are arranged in a fixed pattern within the crystal lattice[1]. This is why the formula for a hydrated compound is written with a dot, followed by a number indicating how many water molecules are attached to one formula unit of the main compound. For example, copper(II) sulfate pentahydrate is written as $CuSO_4 \cdot 5H_2O$. This means that for every one $CuSO_4$ unit, there are five water molecules integrated into the crystal structure.
Common Hydrated Salts in Everyday Life
Hydrated compounds are not just laboratory curiosities; they are part of our daily lives. You might even have some in your bathroom or laundry room.
| Common Name | Chemical Formula | Use and Characteristic |
|---|---|---|
| Epsom Salt | $MgSO_4 \cdot 7H_2O$ | Used in bath salts and as a fertilizer. It is a white, crystalline solid. |
| Blue Vitriol | $CuSO_4 \cdot 5H_2O$ | Bright blue crystals used in agriculture as a fungicide and in school chemistry experiments. |
| Gypsum | $CaSO_4 \cdot 2H_2O$ | Used to make plaster of Paris and drywall for construction. |
| Washing Soda | $Na_2CO_3 \cdot 10H_2O$ | Used as a water softener in laundry detergents. |
The Reversible Nature of Hydration
One of the most fascinating aspects of hydrated compounds is that the process of gaining and losing water is often reversible. This means you can remove the water and, under the right conditions, put it back.
When a hydrated compound is heated gently, it loses its water of crystallisation. The resulting compound is called an anhydrous salt, meaning "without water." A classic classroom experiment demonstrates this with copper(II) sulfate pentahydrate. The beautiful blue crystals turn into a white powder when heated, as the water is driven off.
The reaction is: $CuSO_4 \cdot 5H_2O (s) \xrightarrow{\Delta} CuSO_4 (s) + 5H_2O (g)$
If you then add a few drops of water to the white anhydrous copper(II) sulfate, it will vigorously turn blue again as it re-forms the pentahydrate. This color change makes it a useful test for the presence of water.
From Plaster to Pools: Practical Applications
The properties of hydrated compounds make them incredibly useful. Let's look at two specific examples: plaster of Paris and pool water testing.
Plaster of Paris: This is the common name for calcium sulfate hemihydrate, $(CaSO_4)_2 \cdot H_2O$. It is made by heating gypsum ($CaSO_4 \cdot 2H_2O$) to about $150^\circ C$. When you mix plaster of Paris with water, it undergoes a chemical reaction and re-hydrates, forming a hard, solid mass of interlocking gypsum crystals. This is why it's perfect for making casts, molds, and surgical bandages.
Pool Water Testing: The cobalt chloride test is a simple way to detect moisture. Cobalt(II) chloride is a hydrate that changes color dramatically. The hexahydrate ($CoCl_2 \cdot 6H_2O$) is pink or red, while the anhydrous form ($CoCl_2$) is blue. Strips of paper coated with cobalt chloride can be used to test for humidity or, in a pool context, to check if a surface is dry. The paper is blue when dry and turns pink when wet.
Common Mistakes and Important Questions
A: No, this is a common misconception. The water molecules in a hydrate are part of the solid crystal structure. They are not liquid water filling pores or gaps. They are chemically bound in a specific arrangement, which is why the number of water molecules is fixed and written in the formula.
A: No, not all compounds can form hydrates. The ability to form a stable hydrate depends on the size and charge of the ions in the compound. Ions with a high charge density (a high charge relative to their size) are more likely to attract and hold water molecules in a crystal lattice. Common hydrate formers include sulfates, carbonates, and chlorides of elements like sodium, magnesium, and copper.
A: Efflorescence is when a hydrated salt spontaneously loses its water of crystallisation to the air, becoming a powder. This happens when the vapour pressure[2] of the hydrate is greater than the vapour pressure of the water in the air. Washing soda ($Na_2CO_3 \cdot 10H_2O$) often does this, forming a white dust. Deliquescence is the opposite. A substance absorbs so much water from the air that it dissolves in it, forming a solution. Calcium chloride ($CaCl_2$) is a strong deliquescent salt and is used in dehumidifiers.
Hydrated compounds are a beautiful and practical demonstration of how water interacts with matter on a molecular level. The water of crystallisation is not a passive occupant but an active architect of a crystal's shape, color, and properties. From the blue of copper sulfate to the hardening of plaster, these compounds show us that water can be a fundamental part of a solid's identity. Understanding hydrates opens a window into diverse fields, from geology and materials science to medicine and agriculture, proving that even the most common substance, water, can create extraordinary structures.
Footnote
[1] Crystal Lattice: A crystal lattice is a highly ordered, repeating three-dimensional arrangement of atoms, ions, or molecules that forms a crystal solid.
[2] Vapour Pressure: Vapour pressure is the pressure exerted by a vapour in thermodynamic equilibrium with its condensed phases (solid or liquid) at a given temperature in a closed system. In simpler terms, it is a measure of a substance's tendency to evaporate.