The Mole Concept: Chemical Formula of a Hydrate – Lab Report Assistant Exercise 1: Water of Hydration Data Table 1. Alum Data Mass (8) Object Aluminum Cup (Empty Aluminum Cup +2.00 grams of Alum Aluminum Cup +Alum After 1st Heating Aluminum Cup+ Alum After 2nd Heating Mass of Released H20 Molecular Mass of H2O Moles of Released H2O O.O Questions: A) Calculate the moles of anhydrous (dry) KAI(SO4)2 that were present in the sample. Show all work including units. o.0035 B) Calculate the ratio of moles of H20 to moles of anhydrous KA(SO)2. Show all work including units. Note: Report the ratio to the closest whole number. C) Write the empirical formula for the hydrated KAl(SO4)2, based on your experimental results and answer to Question 2. Show all work including units. Hint: if the ratio of moles of H20 to moles of anhydrous KAl(SO4)2 was 4, then the empirical formula would be: KAI(SO4)2 4H20. D) Describe any visual differences between the hydrated sample and the dried, anhydrous form E) How would the following errors affect the empirical formula for the compound? That is, will these errors cause the calculated number of moles of water in the hydrate to be artificially high or low? a. The student ran out of time and did not do the second heating. Explain how this error will affect the calculation for the number of moles of water in the hydrate? Will the final answer be artificially high or low? How do you know? b. The student recorded the mass of the cup + sample incorrectly and started with 2.20 g of hydrated compound but used 2.00 g in the calculations. Explain how this error will affect the calculation for the number of moles of water in the hydrate? Will the final answer be artificially high or low? How do you know?
The correct answer and explanation is:
Lab Report: The Mole Concept – Chemical Formula of a Hydrate
Exercise 1: Water of Hydration
Answers to Questions
A) Calculate the moles of anhydrous KAl(SO₄)₂:
Given:
- Mass of anhydrous KAl(SO₄)₂ = 0.0035 moles (provided).
Calculation: The moles of anhydrous KAl(SO₄)₂ present in the sample are already determined as 0.0035 moles.
B) Calculate the ratio of moles of H₂O to moles of anhydrous KAl(SO₄)₂:
Given data:
- Mass of released H₂O = Mass of hydrated sample−Mass after 2nd heating\text{Mass of hydrated sample} – \text{Mass after 2nd heating}
- Molecular mass of H₂O = 18.02 g/mol.
Calculate moles of H₂O: Moles of H₂O=Mass of released H₂O18.02.\text{Moles of H₂O} = \frac{\text{Mass of released H₂O}}{18.02}.
Determine the ratio: Ratio of H₂O to anhydrous KAl(SO₄)₂=Moles of H₂OMoles of anhydrous KAl(SO₄)₂.\text{Ratio of H₂O to anhydrous KAl(SO₄)₂} = \frac{\text{Moles of H₂O}}{\text{Moles of anhydrous KAl(SO₄)₂}}.
Round the ratio to the nearest whole number for the empirical formula.
C) Empirical formula for hydrated KAl(SO₄)₂:
If the calculated ratio of H₂O to KAl(SO₄)₂ is, for example, 12, then the formula is: KAl(SO₄)₂ \cdotp 12H₂O.\text{KAl(SO₄)₂ · 12H₂O}.
This represents the number of water molecules per formula unit of KAl(SO₄)₂.
D) Visual differences between hydrated and anhydrous samples:
- Hydrated sample: Appears crystalline, shiny, or translucent due to the bound water.
- Anhydrous sample: Appears powdery, dull, or opaque because of water loss.
E) Effect of Errors on Calculations:
- Skipping the second heating:
- Impact: The sample retains some water, leading to an artificially high mass of “anhydrous” KAl(SO₄)₂. This results in: Fewer moles of water being calculated.\text{Fewer moles of water being calculated.}
- Conclusion: The calculated number of moles of water in the hydrate will be artificially low, as incomplete water loss skews the results.
- Incorrect initial mass of the hydrated sample:
- Impact: Using 2.20 g instead of the actual 2.00 g inflates the initial mass, causing: An overestimation of the water content.\text{An overestimation of the water content.}
- Conclusion: The final number of moles of water in the hydrate will be artificially high, as the mass difference is exaggerated.
Summary Explanation :
Hydrates are compounds with water molecules integrated into their crystal structure. When heated, these molecules are released as vapor, leaving the anhydrous form. The mole ratio of water to anhydrous compound determines the empirical formula of the hydrate. Accurate measurements and heating are critical to the process.
In this experiment, calculating moles requires precise mass data. If the second heating is skipped, some water remains in the sample. This incomplete dehydration increases the apparent mass of the anhydrous compound, leading to fewer moles of water being calculated, artificially lowering the ratio and altering the empirical formula. Similarly, recording an incorrect initial mass inflates the water content calculation, resulting in a higher ratio of water to anhydrous compound, artificially increasing the formula’s water count.
The visual changes in the sample also confirm water loss. Hydrated samples are crystalline and shiny, while anhydrous forms are powdery and dull due to dehydration. By addressing errors and observing proper experimental techniques, students can reliably determine the hydrate’s chemical formula, such as KAl(SO₄)₂·12H₂O.