Problem 157 - Chemistry and Chemical Reactivity

Question

Two general chemistry students working together in the lab weigh out \(0.832 \mathrm{g}\) of \(\mathrm{CaCl}_{2} \cdot 2 \mathrm{H}_{2} \mathrm{O}\) into a crucible. After heating the sample for a short time and allowing the crucible to cool, the students determine that the sample has a mass of \(0.739 \mathrm{g} .\) They then do a quick calculation. On the basis of this calculation, what should they do next? (a) Congratulate themselves on a job well done. (b) Assume the bottle of \(\mathrm{CaCl}_{2} \cdot 2 \mathrm{H}_{2} \mathrm{O}\) was mislabeled; it actually contained something different. (c) Heat the crucible again, and then reweigh it.

Step-by-Step Solution

Answer

The students should heat the crucible again and reweigh it.

Step 1: Identify the Reaction and Mass Loss

The students are heating a hydrate, \(\mathrm{CaCl}_{2} \cdot 2 \mathrm{H}_{2} \mathrm{O}\), which loses water upon heating to form anhydrous \(\mathrm{CaCl}_{2}\). The mass loss should correspond to the total mass of water removed from the hydrate.

Step 2: Calculate Expected Mass Loss

The molar mass of \(\mathrm{CaCl}_{2} \cdot 2 \mathrm{H}_{2} \mathrm{O}\) is approximately 147.02 g/mol. The mass of water lost when the hydrate becomes anhydrous is equal to 2 times the molar mass of water, which is approximately 36.04 g (since the molar mass of water is 18.02 g/mol). Thus, the percentage of water in the hydrate is \((36.04/147.02) \times 100 \approx 24.5\%\).

Step 3: Determine Experimental Mass Loss Percentage

The initial mass of the hydrate was 0.832 g, and the final mass after heating was 0.739 g. The mass lost is \(0.832 \mathrm{g} - 0.739 \mathrm{g} = 0.093 \mathrm{g}\). The experimental mass loss percentage is \((0.093/0.832) \times 100 \approx 11.2\%\).

Step 4: Compare Expected and Experimental Results

The expected mass loss due to water is approximately 24.5%, while the experimental mass loss is about 11.2%. These values significantly differ, indicating that not all water was removed during heating.

Step 5: Final Step: Decide the Action Based on Calculation

Since the experimental mass loss is much lower than expected, the most reasonable step is (c) Heat the crucible again, and then reweigh it. This suggests that the reaction wasn't completed, and more water needs to be driven off.

Key Concepts

Chemical Reactions

Chemical reactions involve the transformation of substances through breaking and forming of bonds. In the context of the lab exercise, the students are dealing with a chemical reaction that occurs when a hydrate, specifically \( \text{CaCl}_2\cdot2\text{H}_2\text{O} \), is heated. A hydrate is a compound that contains water molecules integrated into its crystal structure.
Heating the hydrate causes it to undergo a transformation into an anhydrous form, meaning it releases its water content. This change is a decomposition reaction, where the water molecules separate from the salt compound. Understanding this kind of reaction is key to determining how much the final mass should decrease once the sample is heated thoroughly.
Students must comprehend that the mass loss observed is due to the chemical reaction of water being expelled from the hydrate structure, thus changing its physical properties and mass.

Mass Loss

Mass loss is an essential concept in stoichiometry and chemical reactions, as it indicates a transformation within a substance.
When the students heated the \( \text{CaCl}_2\cdot2\text{H}_2\text{O} \), they observed a decrease in mass, which should represent the water that evaporated from the compound. This process showcases a fundamental application of the law of conservation of mass, illustrating that mass is neither lost nor created, but merely changes form.
In this experiment, knowing the theoretical mass loss from complete dehydration of the hydrate is crucial. The calculated expected mass loss percentage can be contrasted against actual changes observed to confirm if the reaction completed, or if further heating is necessary.

Hydrates

Hydrates are compounds that include water molecules within their crystalline structure. These water molecules, called "water of hydration," are essential to the physical properties of the material. In the exercise, \( \text{CaCl}_2\cdot2\text{H}_2\text{O} \) is used as a hydrate example.
When a hydrate like calcium chloride dihydrate is heated, it undergoes a dehydration process where it releases its water of crystallization. This transforms it into its anhydrous form, in this case, \( \text{CaCl}_2 \). This characteristic loss of water under heating is a definitive trait of hydrates and is utilized to determine purity or composition in lab settings.

Experimental Procedure

An experimental procedure in chemistry often involves several methodical steps to ensure accuracy and reliability of results. In this hydrate experiment, the students followed steps such as:

• Weighing the sample initially and then after heating
• Ensuring the sample is heated adequately to drive off water
• Calculating both the expected and actual mass loss based on theoretical data

Each stage of the experiment requires precision and understanding. When expectations do not match results—as in this case—it often requires revisiting experimental steps.
Analysis of discrepancies is essential, as it can point to incomplete reactions or errors in procedure. Students should measure carefully with analytical balances, maintain consistent heating, and allow the substance to cool properly for accurate measurement. Their solution to heat and weigh again suggests thoroughness in achieving consistent, reliable outcomes in experimental chemistry.