Empirical Formula Of Copper Sulfate Hydrate

Determining the Empirical Formula of Copper Sulfate Hydrate: A Comprehensive Guide

The empirical formula of a compound represents the simplest whole-number ratio of atoms of each element present in the compound. Determining the empirical formula is a fundamental skill in chemistry, and copper sulfate hydrate (CuSO₄·xH₂O) provides an excellent example of how this can be achieved through a series of straightforward experiments. This guide will walk you through the process, explaining the theory, procedure, and calculations involved in finding the value of 'x' – the number of water molecules associated with each copper sulfate molecule.

Understanding Copper Sulfate Hydrate

Copper sulfate hydrate, also known as copper(II) sulfate pentahydrate, is a vibrant blue crystalline compound. The "hydrate" part of the name indicates that water molecules are incorporated into the crystal structure. These water molecules are not simply trapped within the crystal; they are chemically bound to the copper sulfate, influencing its properties. When heated, these water molecules are released, leaving behind anhydrous copper sulfate, a white powder. Our experiment will exploit this property to determine the value of 'x'.

The Experiment: Dehydrating Copper Sulfate Hydrate

This experiment involves carefully heating a known mass of copper sulfate hydrate to drive off the water molecules. By measuring the mass of the anhydrous copper sulfate remaining, we can calculate the mass of water lost and, subsequently, determine the empirical formula.

Materials Required:

• Copper sulfate hydrate (CuSO₄·xH₂O): The sample you'll be analyzing. Ensure it's relatively pure.
• Crucible and lid: A crucible is a small, heat-resistant ceramic container ideal for heating samples. The lid helps to prevent splattering and loss of sample.
• Bunsen burner or hot plate: A heat source for heating the crucible.
• Clay triangle: Supports the crucible over the Bunsen burner or hot plate.
• Ring stand: Provides a stable platform for the clay triangle.
• Balance: An analytical balance is preferable for accurate mass measurements.
• Desiccator (Optional): A desiccator helps to prevent the anhydrous copper sulfate from reabsorbing moisture from the air.
• Spatula or Scoop: For handling the copper sulfate hydrate.
• Heat-resistant gloves: To protect your hands from heat.
• Goggles: To protect your eyes from potential hazards.

Procedure:

1. Weigh the Crucible and Lid: Carefully weigh the clean, dry crucible and its lid using the analytical balance. Record the mass accurately.

2. Add Copper Sulfate Hydrate: Add approximately 2-3 grams of copper sulfate hydrate to the crucible. Record the mass of the crucible, lid, and copper sulfate hydrate.

3. Heat the Sample: Carefully place the crucible (with the lid slightly ajar to allow water vapor to escape) on the clay triangle supported by the ring stand. Heat the crucible gently at first to avoid splattering. Gradually increase the heat until the blue color of the copper sulfate hydrate fades to a white or pale gray, indicating the complete removal of water. This may take some time; be patient and persistent.

4. Cool and Weigh: Remove the crucible from the heat and allow it to cool completely. This is crucial to prevent errors due to heat affecting the weighing. Once cooled, carefully place the crucible and lid in a desiccator (if available) to prevent rehydration. Weigh the crucible, lid, and anhydrous copper sulfate.

5. Repeat Heating (Optional): To ensure complete dehydration, repeat steps 3 and 4 until the mass remains constant between consecutive weighings. This confirms that all the water has been driven off.

Calculations: Determining the Empirical Formula

Once the experiment is complete, the data collected needs to be carefully analyzed to determine the empirical formula. Here's a step-by-step guide to the calculations:

1. Mass of Copper Sulfate Hydrate: Subtract the mass of the crucible and lid (from step 1) from the mass of the crucible, lid, and copper sulfate hydrate (from step 2). This gives you the initial mass of copper sulfate hydrate.

2. Mass of Anhydrous Copper Sulfate: Subtract the mass of the crucible and lid (from step 1) from the mass of the crucible, lid, and anhydrous copper sulfate (from step 4). This gives you the final mass of anhydrous copper sulfate.

3. Mass of Water Lost: Subtract the mass of anhydrous copper sulfate (from step 2) from the mass of copper sulfate hydrate (from step 1). This represents the mass of water driven off during heating.

4. Moles of Anhydrous Copper Sulfate: Divide the mass of anhydrous copper sulfate (from step 2) by the molar mass of anhydrous copper sulfate (CuSO₄ = 159.61 g/mol). This gives you the number of moles of anhydrous copper sulfate.

5. Moles of Water: Divide the mass of water lost (from step 3) by the molar mass of water (H₂O = 18.02 g/mol). This gives you the number of moles of water.

6. Mole Ratio: Divide the number of moles of water (from step 5) by the number of moles of anhydrous copper sulfate (from step 4). This ratio should be a whole number or very close to one. This whole number represents the value of 'x' in the formula CuSO₄·xH₂O.

7. Empirical Formula: Substitute the whole number value of 'x' obtained in step 6 into the formula CuSO₄·xH₂O. This is the empirical formula of the copper sulfate hydrate.

Example Calculation:

Let's assume the following data was obtained from the experiment:

• Mass of crucible and lid: 25.00 g
• Mass of crucible, lid, and copper sulfate hydrate: 28.50 g
• Mass of crucible, lid, and anhydrous copper sulfate: 27.00 g

1. Mass of Copper Sulfate Hydrate: 28.50 g - 25.00 g = 3.50 g

2. Mass of Anhydrous Copper Sulfate: 27.00 g - 25.00 g = 2.00 g

3. Mass of Water Lost: 3.50 g - 2.00 g = 1.50 g

4. Moles of Anhydrous Copper Sulfate: 2.00 g / 159.61 g/mol = 0.0125 mol

5. Moles of Water: 1.50 g / 18.02 g/mol = 0.0832 mol

6. Mole Ratio (x): 0.0832 mol / 0.0125 mol ≈ 6.66 ≈ 7 (rounding to the nearest whole number)

7. Empirical Formula: CuSO₄·7H₂O

Sources of Error and Precautions:

Several factors can affect the accuracy of the experiment:

• Incomplete Dehydration: If the sample is not heated sufficiently, some water molecules may remain, leading to an erroneously low value of 'x'.
• Rehydration: Anhydrous copper sulfate readily absorbs moisture from the air. Using a desiccator helps minimize this.
• Spattering: Vigorous heating can cause some of the sample to splatter out of the crucible, resulting in a lower final mass.
• Impurities in the sample: The presence of impurities in the copper sulfate hydrate will affect the results.

Conclusion:

Determining the empirical formula of copper sulfate hydrate is a valuable experiment that reinforces fundamental concepts in stoichiometry and laboratory techniques. By carefully following the procedure and performing the calculations accurately, you can successfully determine the number of water molecules associated with each copper sulfate molecule and gain a deeper understanding of hydrates. Remember, accuracy and precision are crucial for obtaining reliable results. Careful attention to detail and repetition of the heating steps will improve the accuracy of the experiment. The experiment also offers excellent opportunities to discuss experimental errors and improve laboratory skills.