5. Write a balanced thermochemical equation for the heat of combustion of gaseous ethylene, C2H4, given the following enthalpies of formation (kJ/mol): C2H4 (g), 52.3; H20 (I), -285.8; and CO2 (g), -393.5.
5. Write a balanced thermochemical equation for the heat of combustion of gaseous ethylene, C2H4, given the following enthalpies of formation (kJ/mol): C2H4 (g), 52.3; H20 (I), -285.8; and CO2 (g), -393.5.
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
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Problem 1RQ: Define and explain the differences between the following terms. a. law and theory b. theory and...
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![**Problem 5:**
Write a balanced thermochemical equation for the heat of combustion of gaseous ethylene, \( \text{C}_2\text{H}_4 \), given the following enthalpies of formation (kJ/mol):
- \(\text{C}_2\text{H}_4 (g)\): 52.3
- \(\text{H}_2\text{O} (l)\): -285.8
- \(\text{CO}_2 (g)\): -393.5
**Solution:**
To address this problem, you need to write a balanced chemical equation for the combustion of ethylene, then calculate the heat of the reaction using the provided enthalpies of formation.
**Step 1: Write the balanced chemical equation:**
The combustion of ethylene involves its reaction with oxygen to produce carbon dioxide and water. The balanced equation is:
\[ \text{C}_2\text{H}_4 (g) + 3\text{O}_2 (g) \rightarrow 2\text{CO}_2 (g) + 2\text{H}_2\text{O} (l) \]
**Step 2: Calculate the heat of combustion:**
The heat of the reaction (\(\Delta H_{rxn}\)) can be calculated using the enthalpies of formation (\(\Delta H_f^\circ\)):
\[
\Delta H_{rxn} = \left[ 2 \times \Delta H_f^\circ (\text{CO}_2) + 2 \times \Delta H_f^\circ (\text{H}_2\text{O}) \right] - \left[ \Delta H_f^\circ (\text{C}_2\text{H}_4) \right]
\]
Substituting the given values:
\[
\Delta H_{rxn} = \left[ 2 \times (-393.5) + 2 \times (-285.8) \right] - [52.3]
\]
\[
\Delta H_{rxn} = \left[ -787.0 - 571.6 \right] - 52.3
\]
\[
\Delta H_{rxn} = -1358.6 - 52.3
\]
\[
\Delta H_{rxn} = -1410.9 \, \](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2F026c0823-dc6b-4e42-b4f5-f1056ea7aae1%2F20de38aa-91fe-46b8-bc67-c04590485246%2Fu6n647d_processed.png&w=3840&q=75)
Transcribed Image Text:**Problem 5:**
Write a balanced thermochemical equation for the heat of combustion of gaseous ethylene, \( \text{C}_2\text{H}_4 \), given the following enthalpies of formation (kJ/mol):
- \(\text{C}_2\text{H}_4 (g)\): 52.3
- \(\text{H}_2\text{O} (l)\): -285.8
- \(\text{CO}_2 (g)\): -393.5
**Solution:**
To address this problem, you need to write a balanced chemical equation for the combustion of ethylene, then calculate the heat of the reaction using the provided enthalpies of formation.
**Step 1: Write the balanced chemical equation:**
The combustion of ethylene involves its reaction with oxygen to produce carbon dioxide and water. The balanced equation is:
\[ \text{C}_2\text{H}_4 (g) + 3\text{O}_2 (g) \rightarrow 2\text{CO}_2 (g) + 2\text{H}_2\text{O} (l) \]
**Step 2: Calculate the heat of combustion:**
The heat of the reaction (\(\Delta H_{rxn}\)) can be calculated using the enthalpies of formation (\(\Delta H_f^\circ\)):
\[
\Delta H_{rxn} = \left[ 2 \times \Delta H_f^\circ (\text{CO}_2) + 2 \times \Delta H_f^\circ (\text{H}_2\text{O}) \right] - \left[ \Delta H_f^\circ (\text{C}_2\text{H}_4) \right]
\]
Substituting the given values:
\[
\Delta H_{rxn} = \left[ 2 \times (-393.5) + 2 \times (-285.8) \right] - [52.3]
\]
\[
\Delta H_{rxn} = \left[ -787.0 - 571.6 \right] - 52.3
\]
\[
\Delta H_{rxn} = -1358.6 - 52.3
\]
\[
\Delta H_{rxn} = -1410.9 \, \
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