A hydrate is an ionic compound that incorporates water molecules directly within its crystalline structure. This water, often called the water of hydration, is chemically bound but can be driven off easily by heating. The chemical formula of a hydrate is written as \(\text{Salt} \cdot x\text{H}_2\text{O}\). The objective is to determine the whole-number ratio, ‘x’, of water molecules to one formula unit of the anhydrous salt by measuring mass before and after removing the water.
The Experimental Procedure: Measuring Mass Changes
The process begins by preparing a clean crucible and lid, which is heated strongly and allowed to cool to ensure it is completely dry. This equipment is then weighed precisely to establish the initial mass for accurate subtraction.
Next, a sample of the hydrate is placed into the crucible, and the combined mass of the crucible, lid, and hydrate is measured. The mass of the hydrate alone is calculated by subtracting the initial mass of the empty crucible and lid. This establishes the starting mass of the salt plus its water of hydration.
The sample is then heated, usually gently at first, to vaporize the water molecules bound within the crystal structure. The heating must be controlled to ensure all the water is removed without causing the anhydrous salt itself to decompose.
The heating and cooling cycle must continue until a “constant mass” is achieved. This involves reheating, cooling, and reweighing repeatedly until two consecutive mass readings agree. Achieving constant mass confirms that all water has been driven out. The final mass of the crucible and contents is recorded, and the mass of the anhydrous salt is found by subtracting the mass of the empty crucible.
The Essential Calculations: Determining the Mole Ratio
The calculation involves determining the mass of water lost during heating. This mass is found by subtracting the final mass of the anhydrous salt from the initial mass of the hydrated sample.
Once the masses of both the water and the anhydrous salt are known, they must be converted into moles. The molar mass of water (18.02 g/mol) is used to convert the mass of lost water into moles of water. The molar mass of the anhydrous salt is used to convert the remaining salt mass into moles of the anhydrous compound.
Finding the mole ratio is the final step, which reveals the number of water molecules per unit of the salt. This is achieved by dividing the moles of water by the moles of the anhydrous salt.
This division often results in a number very close to a whole integer, such as \(4.98\) or \(5.02\), which can be safely rounded to the nearest whole number, like five. This final whole number represents the stoichiometric coefficient ‘x’ in the hydrate formula.
Interpreting the Final Formula and Accuracy
The final step is to write the completed chemical formula using the determined whole-number ratio, ‘x’. If the anhydrous salt was copper sulfate (\(\text{CuSO}_4\)) and the calculated ratio ‘x’ was five, the final formula would be written as \(\text{CuSO}_4 \cdot 5\text{H}_2\text{O}\). This formula indicates that five molecules of water are associated with every one formula unit of the salt.
A common source of error is incomplete dehydration, where not all the water is removed. This causes the mass of the anhydrous salt to be artificially high, leading to a calculated mass of lost water that is too low. This results in an ‘x’ value that is smaller than the true ratio.
Conversely, overheating the sample can cause the anhydrous salt itself to decompose or react with the atmosphere. If the salt decomposes, the final mass reading will be lower than expected, leading to an overestimation of lost water and an ‘x’ value that is too high. Precision in weighing prevents atmospheric moisture absorption.