12. A C Considering the following two reactions B AG¹ = -41.5 KJ/mol DAG = +6.5 KJ/mol Which of the followings is correct for the reaction A + C → B + D?
12. A C Considering the following two reactions B AG¹ = -41.5 KJ/mol DAG = +6.5 KJ/mol Which of the followings is correct for the reaction A + C → B + D?
Biochemistry
9th Edition
ISBN:9781319114671
Author:Lubert Stryer, Jeremy M. Berg, John L. Tymoczko, Gregory J. Gatto Jr.
Publisher:Lubert Stryer, Jeremy M. Berg, John L. Tymoczko, Gregory J. Gatto Jr.
Chapter1: Biochemistry: An Evolving Science
Section: Chapter Questions
Problem 1P
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Question
![### Understanding Energetics of Chemical Reactions
#### Considering the following two reactions:
- **Reaction 1:** \( A \rightarrow B \), with a standard Gibbs free energy change \( \Delta G^0 = -41.5 \text{ KJ/mol} \)
- **Reaction 2:** \( C \rightarrow D \), with a standard Gibbs free energy change \( \Delta G^0 = +6.5 \text{ KJ/mol} \)
**Problem Statement:**
Which of the following is correct for the reaction \( A + C \rightarrow B + D \)?
**Options:**
A. The reaction must be thermodynamically favorable under cellular conditions
B. The reaction is able to drive the synthesis of ATP under standard conditions
C. The reaction must be endergonic under cellular conditions
D. The reaction is endergonic under standard conditions
### Explanation:
1. **Thermodynamic Favorability and Gibbs Free Energy:**
- A reaction is thermodynamically favorable (exergonic) if \( \Delta G \) is negative, meaning it releases energy.
- An endergonic reaction has a positive \( \Delta G \), indicating it requires an input of energy.
2. **Combining Reactions:**
- To determine the overall \( \Delta G \) for the coupled reaction \( A + C \rightarrow B + D \), add the \( \Delta G^0 \) of both reactions:
\[
\Delta G^0 (\text{overall}) = \Delta G^0 (A \rightarrow B) + \Delta G^0 (C \rightarrow D) \\
\Delta G^0 (\text{overall}) = -41.5 \text{ KJ/mol} + 6.5 \text{ KJ/mol} = -35.0 \text{ KJ/mol}
\]
3. **Interpreting the Result:**
- The overall reaction has a negative \( \Delta G^0 \) of -35.0 KJ/mol, illustrating that it is exergonic and thus thermodynamically favorable.
### Conclusion:
By evaluating the Gibbs free energy changes, we determine that:
**Option A** is the correct answer.
**A. The reaction must be thermodynamically favorable under cellular conditions** (as evidenced by the overall negative \( \Delta G^0 \)).
Others](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2F83cf1c63-6d62-4d43-b856-80d0ca21ff05%2F784366dc-630c-41a6-92dc-3de91f801ef3%2Fwh5ut7ip_processed.png&w=3840&q=75)
Transcribed Image Text:### Understanding Energetics of Chemical Reactions
#### Considering the following two reactions:
- **Reaction 1:** \( A \rightarrow B \), with a standard Gibbs free energy change \( \Delta G^0 = -41.5 \text{ KJ/mol} \)
- **Reaction 2:** \( C \rightarrow D \), with a standard Gibbs free energy change \( \Delta G^0 = +6.5 \text{ KJ/mol} \)
**Problem Statement:**
Which of the following is correct for the reaction \( A + C \rightarrow B + D \)?
**Options:**
A. The reaction must be thermodynamically favorable under cellular conditions
B. The reaction is able to drive the synthesis of ATP under standard conditions
C. The reaction must be endergonic under cellular conditions
D. The reaction is endergonic under standard conditions
### Explanation:
1. **Thermodynamic Favorability and Gibbs Free Energy:**
- A reaction is thermodynamically favorable (exergonic) if \( \Delta G \) is negative, meaning it releases energy.
- An endergonic reaction has a positive \( \Delta G \), indicating it requires an input of energy.
2. **Combining Reactions:**
- To determine the overall \( \Delta G \) for the coupled reaction \( A + C \rightarrow B + D \), add the \( \Delta G^0 \) of both reactions:
\[
\Delta G^0 (\text{overall}) = \Delta G^0 (A \rightarrow B) + \Delta G^0 (C \rightarrow D) \\
\Delta G^0 (\text{overall}) = -41.5 \text{ KJ/mol} + 6.5 \text{ KJ/mol} = -35.0 \text{ KJ/mol}
\]
3. **Interpreting the Result:**
- The overall reaction has a negative \( \Delta G^0 \) of -35.0 KJ/mol, illustrating that it is exergonic and thus thermodynamically favorable.
### Conclusion:
By evaluating the Gibbs free energy changes, we determine that:
**Option A** is the correct answer.
**A. The reaction must be thermodynamically favorable under cellular conditions** (as evidenced by the overall negative \( \Delta G^0 \)).
Others
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