The ∆G values were calculated for a particular set of conditions known as standard conditions.
The standard free energy change (∆Gº’) of a chemical reaction is the amount of energy released in the conversion of reactants to products under standard conditions.
For biochemical reactions, standard conditions are generally defined as 25°C (298 K), 1 M concentrations of all reactants and products, 1 atm, and pH of 7.0 (the prime mark in ∆Gº’ indicates that pH is included in the definition).
The conditions inside a cell or organism can be very different from these standard conditions, so ∆G values for biological reactions in vivo may vary widely from their standard free energy change (∆Gº’) values. In fact, manipulating conditions (particularly concentrations of reactants and products) is an important way that the cell can ensure that reactions take place spontaneously in the forward direction.
To understand why this is the case, it’s useful to bring up the concept of chemical equilibrium.
As a refresher on chemical equilibrium, let’s imagine that we start a reversible reaction with pure reactants (no product present at all).
At first, the forward reaction will proceed rapidly, as there are lots of reactants that can be converted into products.
The reverse reaction, in contrast, will not take place at all, as there are no products to turn back into reactants.
As product accumulates, however, the reverse reaction will begin to happen more and more often.
This process will continue until the reaction system reaches a balance point, called chemical equilibrium, at which the forward and reverse reactions take place at the same rate. At this point, both reactions continue to occur, but the overall concentrations of products and reactants no longer change. Each reaction has its own unique, characteristic ratio of products to reactants at equilibrium.
When a reaction system is at equilibrium, it is in its lowest-energy state possible (has the least possible free energy).
If a reaction is not at equilibrium, it will move spontaneously towards equilibrium, because this allows it to reach a lower-energy, more stable state.
This may mean a net movement in the forward direction, converting reactants to products, or in the reverse direction, turning products back into reactants.
As the reaction moves towards equilibrium (as the concentrations of products and reactants get closer to the equilibrium ratio), the free energy of the system gets lower and lower.
A reaction that is at equilibrium can no longer do any work, because the free energy of the system is as low as possible.
Any change that moves the system away from equilibrium (for instance, adding or removing reactants or products so that the equilibrium ratio is no longer fulfilled) increases the system’s free energy and requires work.
If a cell were an isolated system, its chemical reactions would reach equilibrium, which would not be a good thing. If a cell’s reaction reached equilibrium, the cell would die because there would be no free energy left to perform the work needed to keep it alive.
Cells stay out of equilibrium by manipulating concentrations of reactants and products to keep their metabolic reactions running in the right direction. For instance:
Curious how this pushing and pulling actually works? Check out the reaction coupling video to learn more!