MEDizzy - Consider the reaction shown below. If [A] = 5.00 mM, [B]

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Energy Metabolism Overview

Consider the reaction shown below. If [A] = 5.00 mM, [B] = 2.50 mM, and [C] = 1.25 mM, what would the concentration of D have to be to allow this to be a favorable reaction under these conditions? A + B ⇌ C + D ΔGo′ = +8.65 kcal/mol

ExplanationRecall, ΔG = ΔGo′ + RTln ([C][D]/[A][B]). In order for the reaction to be favorable, ΔG must be negative. If one solves for the concentration of D required for ΔG = 0, then any concentration lower than the one calculated will be sufficient to allow for a negative ΔG. Thus, when ΔG is set to zero, ΔGo′ = −RT ln ([C][D]/[A][B]). Therefore, 8.65 = −(1.98 × 10−3)(298) ln ([1.25D]/[12.5]). This reduces to −14.66 = ln (D/10), or (10)(e−14.66) = [D] in mM. [D] = 4.3 × 10−6 mM, or 4.3 nM.