Explain Water Of Crystallization. Calculate The Oxygen Percentage In CaSO₄. How To Determine The Formula Of Hydrated Calcium Sulphate, Given A Sample?
Calcium sulphate, a versatile compound, exists in both anhydrous (CaSO₄) and hydrated (CaSO₄.nH₂O) forms, each with distinct properties and applications. This exploration delves into the fascinating world of calcium sulphate, focusing on the concept of water of crystallisation, calculating the percentage of oxygen in the anhydrous form, and determining the formula of a hydrated calcium sulphate sample. Understanding these aspects is crucial in various fields, including chemistry, materials science, and construction.
2.1 Calcium Sulphate: Anhydrous (CaSO₄) vs. Hydrated (CaSO₄.nH₂O)
Calcium sulphate is a common chemical compound with the formula CaSO₄. It exists in nature in various forms, the most common being gypsum (CaSO₄·2H₂O). The key difference between these forms lies in the presence and amount of water molecules incorporated into the crystal structure. Anhydrous calcium sulphate (CaSO₄) is the form without any water molecules, while hydrated calcium sulphate (CaSO₄.nH₂O) contains a specific number (n) of water molecules per formula unit. These water molecules are known as water of crystallization and are chemically bound within the crystal lattice.
The significance of hydration lies in its impact on the physical and chemical properties of the compound. For example, gypsum (CaSO₄·2H₂O), a dihydrate, is a soft mineral widely used in construction. When heated, gypsum loses water molecules, transforming into plaster of Paris (CaSO₄·0.5H₂O), a hemihydrate. This process is reversible to some extent; adding water to plaster of Paris allows it to rehydrate and harden, making it a valuable material for casts and molds. Anhydrous calcium sulphate, on the other hand, has different applications due to its lack of water molecules and distinct properties.
The value of 'n' in CaSO₄.nH₂O determines the specific type of hydrated calcium sulphate. Common hydrates include the dihydrate (CaSO₄·2H₂O, gypsum), the hemihydrate (CaSO₄·0.5H₂O, plaster of Paris), and other less common forms. Each hydrate exhibits unique characteristics regarding its stability, solubility, and reactivity. Understanding the hydration state of calcium sulphate is crucial for its effective use in various applications. For instance, the controlled hydration and dehydration of calcium sulphate are essential processes in the production of cement and other construction materials.
2.1.1 Defining Water of Crystallisation
Water of crystallisation, also known as water of hydration, refers to the water molecules that are chemically bound within the crystal lattice of a compound. These water molecules are present in a definite stoichiometric ratio, meaning that a specific number of water molecules are associated with each formula unit of the compound. The presence of water of crystallisation significantly affects the crystal structure and properties of the compound, influencing its stability, appearance, and reactivity.
To further elaborate on the definition of water of crystallisation, it's important to understand how these water molecules are incorporated into the crystal structure. During the crystallisation process, as ions or molecules come together to form a solid, water molecules can become trapped within the lattice. These water molecules are not simply adsorbed onto the surface; they are held in specific positions within the crystal structure by weak chemical bonds, such as hydrogen bonds. This integration of water molecules is crucial for the overall stability and shape of the crystal.
The number of water molecules associated with each formula unit is represented by the 'n' in the chemical formula CaSO₄.nH₂O. For example, in gypsum (CaSO₄·2H₂O), n = 2, indicating that two water molecules are associated with each calcium sulphate unit. This stoichiometric relationship is fundamental to understanding the composition and behaviour of hydrated compounds. Heating a hydrated compound can often remove the water of crystallisation, resulting in a change in the compound's structure and properties. This process is commonly used to prepare anhydrous forms of compounds or to create materials with different characteristics.
Understanding the concept is crucial in various chemical and industrial applications. For instance, in the pharmaceutical industry, the hydration state of a drug can affect its solubility and bioavailability. In the construction industry, the hydration of cement is a critical process that determines its strength and setting time. Therefore, accurately determining and controlling the water of crystallisation is essential for achieving desired material properties and performance.
2.1.2 Calculating the Percentage of Oxygen in Anhydrous Calcium Sulphate, CaSO₄
To calculate the percentage of oxygen in anhydrous calcium sulphate (CaSO₄), we need to determine the molar mass of the compound and the contribution of oxygen to that mass. This calculation involves using the atomic masses of each element present in the compound, which can be found on the periodic table. The process is a fundamental application of stoichiometry, a core concept in chemistry.
First, let's break down the calculation step by step. The atomic masses are as follows: Calcium (Ca) has an atomic mass of approximately 40.08 g/mol, Sulphur (S) has an atomic mass of approximately 32.07 g/mol, and Oxygen (O) has an atomic mass of approximately 16.00 g/mol. In CaSO₄, there is one calcium atom, one sulphur atom, and four oxygen atoms. Therefore, the molar mass of CaSO₄ can be calculated as follows:
Molar mass of CaSO₄ = (1 × Atomic mass of Ca) + (1 × Atomic mass of S) + (4 × Atomic mass of O)
= (1 × 40.08 g/mol) + (1 × 32.07 g/mol) + (4 × 16.00 g/mol)
= 40.08 g/mol + 32.07 g/mol + 64.00 g/mol
= 136.15 g/mol
Now that we have the molar mass of CaSO₄, we can calculate the mass contribution of oxygen. There are four oxygen atoms in the compound, each with an atomic mass of 16.00 g/mol, so the total mass of oxygen is 4 × 16.00 g/mol = 64.00 g/mol. The percentage of oxygen in CaSO₄ is then calculated as:
Percentage of Oxygen = (Mass of Oxygen / Molar mass of CaSO₄) × 100%
= (64.00 g/mol / 136.15 g/mol) × 100%
≈ 47.01%
Therefore, the percentage of oxygen in anhydrous calcium sulphate (CaSO₄) is approximately 47.01%. This calculation demonstrates the importance of stoichiometry in determining the composition of chemical compounds. The percentage composition of elements in a compound is crucial for understanding its properties and reactivity.
2.1.3 Determining the Formula of Hydrated Calcium Sulphate
To determine the formula of hydrated calcium sulphate (CaSO₄.nH₂O), an experimental approach is necessary. This typically involves heating a known mass of the hydrated salt to drive off the water of crystallisation and then measuring the mass of the remaining anhydrous salt. By comparing the masses of the hydrated and anhydrous forms, we can calculate the number of water molecules (n) associated with each formula unit of CaSO₄.
The experimental procedure generally follows these steps: First, a known mass of the hydrated calcium sulphate is accurately weighed. This initial mass represents the combined mass of CaSO₄ and the water of crystallisation. The sample is then heated strongly to evaporate the water molecules. The heating process should be carried out until the mass of the sample remains constant, indicating that all the water has been driven off. The anhydrous calcium sulphate is then cooled in a desiccator to prevent the absorption of moisture from the air, and its mass is measured.
From the experimental data, we can calculate the mass of water lost during heating by subtracting the mass of the anhydrous salt from the mass of the hydrated salt. This mass of water is then converted to moles using the molar mass of water (18.015 g/mol). Similarly, the mass of the anhydrous CaSO₄ is converted to moles using its molar mass (136.15 g/mol). The ratio of moles of water to moles of CaSO₄ gives us the value of 'n', which represents the number of water molecules per formula unit in the hydrated salt.
For example, let's say we start with 5.00 g of a hydrated calcium sulphate and, after heating, we obtain 3.92 g of anhydrous CaSO₄. The mass of water lost is 5.00 g - 3.92 g = 1.08 g. Converting these masses to moles:
Moles of H₂O = 1.08 g / 18.015 g/mol ≈ 0.06 mol Moles of CaSO₄ = 3.92 g / 136.15 g/mol ≈ 0.029 mol
Now, we calculate the ratio:
n = Moles of H₂O / Moles of CaSO₄ = 0.06 mol / 0.029 mol ≈ 2
Therefore, in this example, the formula of the hydrated calcium sulphate is CaSO₄·2H₂O, which is gypsum. This experimental determination of the formula highlights the practical application of stoichiometry in identifying chemical compounds and understanding their composition. The accuracy of the results depends on precise measurements and careful execution of the experimental procedure.
In summary, calcium sulphate's existence in both anhydrous and hydrated forms showcases the significance of water of crystallisation in determining a compound's properties. The calculation of the percentage of oxygen in anhydrous calcium sulphate demonstrates the practical application of stoichiometry. Furthermore, the experimental determination of the formula for hydrated calcium sulphate underscores the importance of empirical methods in chemistry. Understanding these concepts is crucial for various scientific and industrial applications, making the study of calcium sulphate a valuable exercise in chemistry.