Which of the following reactions are spontaneous (favorable)?
A. DHAP ---> glyceraldehyde - 3 - phosphate ΔG=3.8 kJ/mol
B. C2H4+H2 ----> C2H6 ΔG=-150.97kJ/mol
C. C4H4O5 ---> C4H2O4+H2O ΔG=3.1 kJ/mol
D. L-malata + NAD + ---> oxalocacetate + NADH+H^+ ΔG=29.7 kJ/mol
E. glutamate + NAD^+ H2O ---> NH4^+ a-ketoglutarate+NADH+H^+ ΔG=3.7 kcal/mol
F. C6H13O9P+ATP ---> C6H14O12P2+ADP ΔG=-14.2 kJ/mol
To be spontaneous dG must be - (that's negative).
To determine whether a reaction is spontaneous or favorable, you need to look at the sign of the Gibbs free energy change (ΔG). If ΔG is negative, the reaction is spontaneous or favorable. If ΔG is positive, the reaction is non-spontaneous or unfavorable.
Let's go through the given reactions one by one:
A. DHAP ---> glyceraldehyde - 3 - phosphate ΔG = 3.8 kJ/mol
Since ΔG is positive (3.8 kJ/mol), Reaction A is non-spontaneous or unfavorable.
B. C2H4 + H2 ----> C2H6 ΔG = -150.97 kJ/mol
Since ΔG is negative (-150.97 kJ/mol), Reaction B is spontaneous or favorable.
C. C4H4O5 ---> C4H2O4 + H2O ΔG = 3.1 kJ/mol
Since ΔG is positive (3.1 kJ/mol), Reaction C is non-spontaneous or unfavorable.
D. L-malate + NAD+ ---> oxaloacetate + NADH+H+ ΔG = 29.7 kJ/mol
Since ΔG is positive (29.7 kJ/mol), Reaction D is non-spontaneous or unfavorable.
E. glutamate + NAD+ + H2O ---> NH4+ + α-ketoglutarate + NADH+H+ ΔG = 3.7 kcal/mol
The given ΔG is in kcal/mol, so we need to convert it to kJ/mol by multiplying it by 4.184. In this case, ΔG = 3.7 kcal/mol * 4.184 = 15.4 kJ/mol. Since ΔG is positive (15.4 kJ/mol), Reaction E is non-spontaneous or unfavorable.
F. C6H13O9P + ATP ---> C6H14O12P2 + ADP ΔG = -14.2 kJ/mol
Since ΔG is negative (-14.2 kJ/mol), Reaction F is spontaneous or favorable.
In summary, the reactions that are spontaneous (favorable) are:
B. C2H4 + H2 ----> C2H6 ΔG = -150.97 kJ/mol
F. C6H13O9P + ATP ---> C6H14O12P2 + ADP ΔG = -14.2 kJ/mol