Molecules A and B contain 110 kcal/mol of free energy, and molecules B and C contain 150 kcal/mol of energy. A and B are converted to C and D. What can be concluded?

• A. A and B will be converted to C and D with a net release of energy.
• B. The conversion of A and B to C and D is exergonic; the products are less organized than the reactants.
• C. The conversion of A and B to C and D is spontaneous.
• D. The entropy in the products, C and D, is higher than in the reactants, A and B.
• E. The reaction that proceeds to convert A and B to C and D is endergonic; the products are more organized than the reactants.

Answer: (E)

Verified step by step guidance

1. Begin by understanding the concept of free energy in chemical reactions. Free energy, often denoted as Gibbs free energy (G), is the energy available to do work. In a chemical reaction, the change in free energy (ΔG) determines whether the reaction is spontaneous or requires energy input.
2. Calculate the total free energy of the reactants (A and B) and the products (C and D). Molecules A and B together contain 110 kcal/mol of free energy, while molecules B and C contain 150 kcal/mol. Since B is common in both pairs, focus on the net change in energy.
3. Determine the change in free energy (ΔG) for the reaction. The formula for ΔG is: ΔG = G_products - G_reactants. Substitute the values: ΔG = (energy of C and D) - (energy of A and B).
4. Analyze the sign of ΔG. If ΔG is positive, the reaction is endergonic, meaning it requires energy input and the products are more organized than the reactants. If ΔG is negative, the reaction is exergonic, meaning it releases energy and the products are less organized.
5. Conclude based on the calculated ΔG. Since the problem states that the reaction is endergonic, it implies that the products (C and D) are more organized than the reactants (A and B), and the reaction requires an input of energy.