Assume that 0.321 g of diborane is combusted in a calorimeter whose heat capacity (Calorimeter) is 7.854 kJ/°C at 16.31°C. What is the final temperature of the calorimeter? °C 1st attempt
Assume that 0.321 g of diborane is combusted in a calorimeter whose heat capacity (Calorimeter) is 7.854 kJ/°C at 16.31°C. What is the final temperature of the calorimeter? °C 1st attempt
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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![**Text for Educational Website:**
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**Problem Statement:**
Assume that 0.321 g of diborane is combusted in a calorimeter whose heat capacity (\(C_{\text{calorimeter}}\)) is 7.854 kJ/°C at 16.31°C. What is the final temperature of the calorimeter?
\[
\text{Final Temperature:} \ \_\_\_\_ \ °C
\]
**Instructions for Solution Attempt:**
1. **Calculate the Heat Released:**
- Use the molecular weight of diborane and its combustion enthalpy to determine the heat released upon combustion.
2. **Apply Calorimetry Principle:**
- Use the formula \( q = C \times \Delta T \), where \( q \) is the heat absorbed by the calorimeter, \( C \) is the calorimeter's heat capacity, and \( \Delta T \) is the temperature change.
3. **Determine the Final Temperature:**
- Rearrange the equation to solve for the final temperature.
**First Attempt:**
- Complete the calculation using the steps above and compare with experimental data if available.
---
**Graph or Diagram:**
- *There is no graph or diagram provided in the image.*
This structured approach will help students understand how to approach calorimetry problems and apply thermodynamic principles in a systematic manner.](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2Fc2367310-44c7-43e5-b110-21237a7a5b46%2F19946a3f-20d8-4ca0-9e5d-be3c0fb85c04%2Fadtfjnx_processed.jpeg&w=3840&q=75)
Transcribed Image Text:**Text for Educational Website:**
---
**Problem Statement:**
Assume that 0.321 g of diborane is combusted in a calorimeter whose heat capacity (\(C_{\text{calorimeter}}\)) is 7.854 kJ/°C at 16.31°C. What is the final temperature of the calorimeter?
\[
\text{Final Temperature:} \ \_\_\_\_ \ °C
\]
**Instructions for Solution Attempt:**
1. **Calculate the Heat Released:**
- Use the molecular weight of diborane and its combustion enthalpy to determine the heat released upon combustion.
2. **Apply Calorimetry Principle:**
- Use the formula \( q = C \times \Delta T \), where \( q \) is the heat absorbed by the calorimeter, \( C \) is the calorimeter's heat capacity, and \( \Delta T \) is the temperature change.
3. **Determine the Final Temperature:**
- Rearrange the equation to solve for the final temperature.
**First Attempt:**
- Complete the calculation using the steps above and compare with experimental data if available.
---
**Graph or Diagram:**
- *There is no graph or diagram provided in the image.*
This structured approach will help students understand how to approach calorimetry problems and apply thermodynamic principles in a systematic manner.
![A bomb calorimeter can be used to measure the enthalpy of reaction for combustions, such as the combustion of diborane represented by the following balanced equation:
\[ \text{B}_2\text{H}_6 (g) + 3\text{O}_2 (g) \rightarrow \text{B}_2\text{O}_3 (s) + 3\text{H}_2\text{O} (g) \]
\(\Delta H_{\text{combustion}}\) for this process is \(-1958 \, \text{kJ/mol}\).](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2Fc2367310-44c7-43e5-b110-21237a7a5b46%2F19946a3f-20d8-4ca0-9e5d-be3c0fb85c04%2Fdat2fku_processed.jpeg&w=3840&q=75)
Transcribed Image Text:A bomb calorimeter can be used to measure the enthalpy of reaction for combustions, such as the combustion of diborane represented by the following balanced equation:
\[ \text{B}_2\text{H}_6 (g) + 3\text{O}_2 (g) \rightarrow \text{B}_2\text{O}_3 (s) + 3\text{H}_2\text{O} (g) \]
\(\Delta H_{\text{combustion}}\) for this process is \(-1958 \, \text{kJ/mol}\).
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