Remember that he actually ended up with 2.9 g of CaO. This is a decomposition reaction which involves heating up the reactant in order to break it down and release a gas, so how would you explain that his actual yield was higher than his theoretical yield?
Remember that he actually ended up with 2.9 g of CaO. This is a decomposition reaction which involves heating up the reactant in order to break it down and release a gas, so how would you explain that his actual yield was higher than his theoretical yield?
Chemistry
10th Edition
ISBN:9781305957404
Author:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste
Publisher:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste
Chapter1: Chemical Foundations
Section: Chapter Questions
Problem 1RQ: Define and explain the differences between the following terms. a. law and theory b. theory and...
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Remember that he actually ended up with 2.9 g of CaO. This is a decomposition reaction which involves heating up the reactant in order to break it down and release a gas, so how would you explain that his actual yield was higher than his theoretical yield?
![### Understanding the Decomposition of Calcium Carbonate
**Read the passage and use the equation to answer the rest of the questions.**
The molecule CaCO₃ is what makes up eggshells, and it can decompose (break down) when heated strongly. This decomposition reaction is given by the equation:
\[ \text{CaCO}_3(\text{s}) \rightarrow \text{CaO}(\text{s}) + \text{CO}_2(\text{g}) \]
Your friend argues that the chemical equation doesn't follow the law of conservation of mass. He says, "Moles are not conserved since 1 mole of reactants produced 2 moles of product. This means the mass is higher in the end." You think he's wrong, and want to explain and do an experiment to prove it.
---
**Explaining the Experiment to Prove the Law of Conservation of Mass:**
1. **Gather Materials:**
- A sample of calcium carbonate (CaCO₃)
- A heat source to decompose the CaCO₃
- Balance to measure mass
- A container to collect the products formed
2. **Procedure:**
- Measure the initial mass of the calcium carbonate sample.
- Heat the sample strongly to decompose it into calcium oxide (CaO) and carbon dioxide (CO₂).
- Measure the mass of the calcium oxide residue.
- Collect and measure the mass of carbon dioxide released (either directly or using an indirect method like volume).
3. **Analysis:**
- Compare the initial mass of CaCO₃ with the combined mass of CaO and CO₂.
- According to the law of conservation of mass, the total mass before and after the reaction should be equal.
- The chemical equation shows that 1 mole of CaCO₃ decomposes into 1 mole of CaO and 1 mole of CO₂, proving mass conservation in terms of molecular weight.
This experiment will demonstrate that while the number of moles may look different, the total mass remains the same, thus validating the law of conservation of mass.
---
**Graphical Representation:**
While no graphs or diagrams were included in the text, a helpful diagram to support this explanation might include:
- A balanced chemical equation showing molecular weights.
- A before-and-after mass comparison table or bar graph.
- A visual setup of the experiment showing the heating process and collection of products.](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2F46ade868-a4ca-4dcd-b69f-bba1f121fc8f%2F90bfbc6e-7cce-4f58-b5c0-1fd3334f36d4%2Fx15xtej_processed.jpeg&w=3840&q=75)
Transcribed Image Text:### Understanding the Decomposition of Calcium Carbonate
**Read the passage and use the equation to answer the rest of the questions.**
The molecule CaCO₃ is what makes up eggshells, and it can decompose (break down) when heated strongly. This decomposition reaction is given by the equation:
\[ \text{CaCO}_3(\text{s}) \rightarrow \text{CaO}(\text{s}) + \text{CO}_2(\text{g}) \]
Your friend argues that the chemical equation doesn't follow the law of conservation of mass. He says, "Moles are not conserved since 1 mole of reactants produced 2 moles of product. This means the mass is higher in the end." You think he's wrong, and want to explain and do an experiment to prove it.
---
**Explaining the Experiment to Prove the Law of Conservation of Mass:**
1. **Gather Materials:**
- A sample of calcium carbonate (CaCO₃)
- A heat source to decompose the CaCO₃
- Balance to measure mass
- A container to collect the products formed
2. **Procedure:**
- Measure the initial mass of the calcium carbonate sample.
- Heat the sample strongly to decompose it into calcium oxide (CaO) and carbon dioxide (CO₂).
- Measure the mass of the calcium oxide residue.
- Collect and measure the mass of carbon dioxide released (either directly or using an indirect method like volume).
3. **Analysis:**
- Compare the initial mass of CaCO₃ with the combined mass of CaO and CO₂.
- According to the law of conservation of mass, the total mass before and after the reaction should be equal.
- The chemical equation shows that 1 mole of CaCO₃ decomposes into 1 mole of CaO and 1 mole of CO₂, proving mass conservation in terms of molecular weight.
This experiment will demonstrate that while the number of moles may look different, the total mass remains the same, thus validating the law of conservation of mass.
---
**Graphical Representation:**
While no graphs or diagrams were included in the text, a helpful diagram to support this explanation might include:
- A balanced chemical equation showing molecular weights.
- A before-and-after mass comparison table or bar graph.
- A visual setup of the experiment showing the heating process and collection of products.
Transcribed Image Text:**BONUS!** Remember that he actually ended up with 2.9 g of CaO. This is a decomposition reaction which involves heating up the reactant in order to break it down & release a gas, so how would you explain that his actual yield was *higher* than his theoretical yield?
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