Question

The reaction catalyzed by malate dehydrogenase,

$Malate=NA{D}^{+}\to oxalocetate+NAD+H^{+}$

has a $∆G^{o^{\text{'}}}$ value of $+29.7kJ.mo{l}^{-1}$ . (a) Would this reaction occur spontaneously in a cell? (b) How does the citrate synthase reaction (described in Problem 7 ) promote the malate dehydrogenase reaction in the cell? What is the overall change in free energy for the two reactions?

Short answer

⦁ In a cell, the reaction would not happen on its own.

⦁ Citrate synthase produces energy and consumes the first reaction's product.

Step-by-step solution

Synthesis Reaction

When several reactants unite to create a single product, the reaction is referred to as a synthesis reaction. Synthesis reactions are exothermic because they release energy in the form of heat and light.

Explanation for part (a)

The formation of oxaloacetate from malate is an endergonic process as the reaction's $∆G^{o^{\text{'}}}$ is positive, which necessitates energy input. Hence, the reaction does not occur spontaneously in the cell.

Explanation for part (b)

By employing it as a reactant, the exergonic citrate synthase reaction supplies the energy necessary for the production of oxaloacetate, which according to LeChatelier's principle propels the process ahead.

The net change in free energy is then seen as:

$∆G^{o^{\text{'}}}=29.7kJ/mol-31.5kJ/mol$

$=1.8kJ/mol$

Hence, the net change is $1.8kJ/mol$.