The ΔG°' for the reaction Citrate isocitrate is +6.64 kJmol-1= +1.59 kcalmol-1. The ΔG°' for the reaction Isocitrate α-Ketoglutarate is -267kJmol-1 = -63.9 kcalmol-1 What is the ΔG°' for the conversion of citrate to α-Ketoglutarate? Is that reaction exergonic or endergonic, and why?
Please explain this to me I have no idea how to solve this.
it's exergonic, I think because it's breaking down the steps from citrate to alpha-ketoglutarate (being the "first" step). when bonds break, it's releasing energy.
To find the ΔG°' for the conversion of citrate to α-Ketoglutarate, we need to use the equation:
ΔG°' = ΔG°' (Isocitrate α-Ketoglutarate) - ΔG°' (Citrate isocitrate)
Given:
ΔG°' (Citrate isocitrate) = +6.64 kJ/mol (or +1.59 kcal/mol)
ΔG°' (Isocitrate α-Ketoglutarate) = -267 kJ/mol (or -63.9 kcal/mol)
Substituting these values into the equation, we get:
ΔG°' = -267 kJ/mol - (+6.64 kJ/mol)
ΔG°' = -267 kJ/mol - 6.64 kJ/mol
ΔG°' = -273.64 kJ/mol
To determine whether the reaction is exergonic (spontaneous) or endergonic (non-spontaneous), we look at the sign of ΔG°'.
Since ΔG°' is negative (-273.64 kJ/mol), the reaction of converting citrate to α-Ketoglutarate is exergonic. This means that the reaction is spontaneous and releases energy.
To find the ΔG°' for the conversion of citrate to α-ketoglutarate, we can use the concept of ΔG°' being additive for a series of reactions.
First, let's understand the meaning of ΔG°'. ΔG°' represents the standard Gibbs free energy change of a reaction under standard conditions, which typically include a temperature of 25°C, a pressure of 1 atm, and all reactants and products at standard concentrations.
In this case, we have two reactions:
1. Citrate → Isocitrate
2. Isocitrate → α-Ketoglutarate
To find the ΔG°' for the conversion of citrate to α-ketoglutarate, we need to add the ΔG°' values for these two reactions.
ΔG°' (Citrate → Isocitrate) = +6.64 kJmol⁻¹ = +1.59 kcalmol⁻¹
ΔG°' (Isocitrate → α-Ketoglutarate) = -267 kJmol⁻¹ = -63.9 kcalmol⁻¹
Now, to find the ΔG°' for the overall reaction, add these two values:
ΔG°' (Citrate → α-Ketoglutarate) = ΔG°' (Citrate → Isocitrate) + ΔG°' (Isocitrate → α-Ketoglutarate)
ΔG°' (Citrate → α-Ketoglutarate) = +1.59 kcalmol⁻¹ + (-63.9 kcalmol⁻¹)
= -62.31 kcalmol⁻¹ (approximately)
The ΔG°' value is negative, indicating an exergonic reaction. An exergonic reaction releases energy. In this case, the conversion of citrate to α-ketoglutarate is exergonic because the overall ΔG°' is negative, meaning that the products have lower free energy than the reactants.