Percent Water In A Hydrate Lab Answer Key

Determining the Percent Water in a Hydrate: A Comprehensive Lab Guide

Determining the percentage of water in a hydrate is a fundamental experiment in chemistry, teaching students about stoichiometry, hydrates, and experimental techniques. This comprehensive guide will walk you through the process, explain the underlying chemistry, provide troubleshooting tips, and offer sample calculations to help you master this essential lab.

Understanding Hydrates

Before diving into the experiment, let's clarify what hydrates are. Hydrates are crystalline compounds that contain water molecules within their crystal structure. This water is chemically bound, not simply absorbed. The water molecules are incorporated in a specific ratio with the anhydrous salt (the salt without water). This ratio is represented in the chemical formula, for example, copper(II) sulfate pentahydrate is written as CuSO₄·5H₂O, indicating five water molecules per formula unit of copper(II) sulfate.

The water molecules in hydrates are often weakly bonded and can be removed by heating, a process called dehydration. This dehydration process is the basis of our experiment to determine the percentage of water in a hydrate.

The Experiment: Determining Percent Water in a Hydrate

This experiment involves heating a known mass of a hydrate until all the water is driven off. By measuring the mass of the remaining anhydrous salt, we can calculate the mass of water lost and then determine the percentage of water in the original hydrate.

Materials:

• Crucible and lid
• Clay triangle
• Ring stand and iron ring
• Bunsen burner or hot plate
• Hydrate sample (e.g., copper(II) sulfate pentahydrate, Epsom salt)
• Analytical balance
• Desiccator (optional, for cooling in a dry environment)
• Tongs

Procedure:

1. Weigh the crucible and lid: Carefully weigh the clean, dry crucible and its lid using an analytical balance. Record this mass (m₁).

2. Add the hydrate: Add a sample of the hydrate to the crucible. Aim for a sample size of approximately 2-3 grams. Record the mass of the crucible, lid, and hydrate (m₂).

3. Heat the hydrate: Place the crucible and lid on the clay triangle supported by the ring stand. Carefully heat the crucible using a Bunsen burner or hot plate, gently at first to avoid splattering. Heat strongly for 10-15 minutes, ensuring the bottom of the crucible is heated evenly. The hydrate will likely change color as the water is driven off.

4. Cool and weigh: Allow the crucible and its contents to cool completely to room temperature. A desiccator can be used to prevent the anhydrous salt from absorbing atmospheric moisture. Once cooled, weigh the crucible, lid, and anhydrous salt (m₃).

5. Repeat steps 3 and 4: To ensure accuracy, repeat steps 3 and 4 until a constant mass is achieved. This means the mass of the crucible and anhydrous salt does not change significantly between consecutive weighings. This ensures all the water has been removed.

Calculations: Determining Percent Water

The percentage of water in the hydrate is calculated using the following formula:

% Water = [(m₂ - m₃) / (m₂ - m₁)] x 100%

Where:

• m₁ = mass of the empty crucible and lid (g)
• m₂ = mass of the crucible, lid, and hydrate (g)
• m₃ = mass of the crucible, lid, and anhydrous salt (g)

This calculation is based on the principle that the difference between m₂ and m₃ represents the mass of water lost during heating, and the difference between m₂ and m₁ represents the mass of the hydrate sample.

Example Calculation:

Let's say:

• m₁ = 25.00 g
• m₂ = 28.50 g
• m₃ = 27.20 g

Then:

% Water = [(28.50 g - 27.20 g) / (28.50 g - 25.00 g)] x 100% = (1.30 g / 3.50 g) x 100% = 37.1%

Therefore, the hydrate sample contains approximately 37.1% water by mass.

Error Analysis and Troubleshooting

Several factors can contribute to errors in this experiment:

• Incomplete dehydration: If the hydrate is not heated sufficiently, some water may remain, leading to an underestimate of the percentage of water. Ensure consistent and thorough heating.

• Spattering: Rapid heating can cause the hydrate to splatter, leading to a loss of sample and inaccurate results. Start with gentle heating and gradually increase the temperature.

• Absorption of atmospheric moisture: The anhydrous salt can absorb moisture from the air after cooling. Use a desiccator to minimize this.

• Impurities in the hydrate: The presence of impurities in the hydrate sample can affect the results. Use a pure hydrate sample whenever possible.

• Inaccurate weighing: Inaccurate measurements using the balance will lead to errors in the final calculation. Ensure the balance is properly calibrated and use appropriate weighing techniques.

Advanced Considerations: Determining the Formula of a Hydrate

Once you have determined the percentage of water in the hydrate, you can use this information, along with the molar mass of the anhydrous salt and water, to determine the empirical formula of the hydrate, which shows the ratio of anhydrous salt to water molecules.

Example:

Let's assume we determined that a hydrate of copper(II) sulfate contains 36.1% water. The molar mass of anhydrous copper(II) sulfate (CuSO₄) is approximately 159.61 g/mol, and the molar mass of water (H₂O) is 18.02 g/mol.

1. Assume 100g of hydrate: If we have 100g of the hydrate, 36.1g is water, and 63.9g is anhydrous CuSO₄.

2. Convert grams to moles:

   • Moles of water = 36.1 g / 18.02 g/mol ≈ 2.00 moles
   • Moles of CuSO₄ = 63.9 g / 159.61 g/mol ≈ 0.40 moles

3. Determine the mole ratio: Divide the number of moles of each component by the smallest number of moles (0.40 moles):

   • Mole ratio of water: 2.00 moles / 0.40 moles ≈ 5
   • Mole ratio of CuSO₄: 0.40 moles / 0.40 moles ≈ 1

4. Write the formula: The empirical formula of the hydrate is CuSO₄·5H₂O, indicating copper(II) sulfate pentahydrate.

This detailed explanation provides a solid foundation for understanding the experiment and performing the calculations accurately. Remember to always prioritize safety precautions when working in a laboratory setting. Thorough understanding of the procedure and meticulous attention to detail will ensure you obtain reliable and meaningful results. By mastering this experiment, you build a strong foundation in experimental chemistry and stoichiometry, skills crucial for further advancements in your chemical studies.