- it is an exergonic reaction itself, and
- it couples itself to the endergonic reaction, so that the overall reaction is exergonic.
The body's ingenious solution is to use another exergonic process to drive the non-spontaneous (endergonic) one. More accurately, it couples the exergonic reaction with the endogonic one so that the free energy released from the former reaction propels the latter one.
For its chosen exergonic reaction, the body generally uses the hydrolysis of ATP (adenosine triphosphate). This is the major 'energy' molecule produced by metabolism, and it serves as a sort of 'energy shuttle': ATP is dispatched to wherever an endergonic reaction needs to take place, and the two reactions are coupled so that the overall reaction is thermodynamically favoured.
To make this more concrete, let's look at the synthesis of Glutamine. This molecule is produced in the following manner:
Glutamic acid + Ammonia → Glutamine
Unfortunately, this reaction is non-spontaneous (ΔG = +3.4 kcal/mol). However, ATP is at hand. The hydrolysis of ATP (ATP + H2O → ADP + Pi) is highly exergonic - (ΔG = -7.3). As you can see if these reactions could somehow be coupled, together the overall reaction would be favoured: ΔG would be -3.9. And so, with the help of ATP, the body can synthesise glutamine without breaking any thermodynamic laws.
But to return to our original question, how exactly is this coupling achieved? The basic answer is this: when ATP is hydrolysed, the free phosphate group (Pi) isn't wasted. Instead, it is transferred to some other molecule, such as the reactant. The molecule is thus phosphorylated. However, this phosphorylated molecule is thereby rendered unstable - it is 'keen' to react all of a sudden. And this new-found tendency allows the desired reaction to take place.
To help make this clearer, I've drawn a diagram of the Glutamic acid + Ammonia → Glutamine reaction above. The big blue molecule is glutamic acid; the other players in the drama are labelled.
In the first stage, the reactants are in place, but since the reaction is non-spontaneous, they are powerless to get going. ATP arrives to help out.
In the second stage, ATP is hydrolysed to ADP + Pi. However, the phosphate group binds to the glutamic acid molecule (phosphorylates it). This phosphorylated intermediate is more reactive than the glutamic acid was by itself.
In the final stage, the ammonia group displaces the phosphate one. This reaction is spontaneous, thanks to the phosphate group making the complex in stage 2 more unstable. Thus, ATP allows glutamine to be formed via two spontaneous reactions, but the price to be paid is that the ATP molecule has been used up - it has been hydrolysed to ADP and Pi.
There are many other examples of ATP's action, but in almost all cases a substance is phosphorylated. As always, this phosphorylated intermediate is more reactive than the original reactant, and so what was once an endergonic reaction is thus rendered exeronic by the coupling of the original reaction with the hydrolysis of ATP.
Source: "Biology", 7th Edn., Campbell & Reece
Good post! Thanks.
ReplyDeletemy teacher used this site :)
ReplyDeleteexplained better than he did!
Excellent, simple, straightforward explanation. I've read about it many times before, but I never came across the explanation of how the phosphate makes the intermediate product less stable. Thanks!
ReplyDeleteI have a question.
ReplyDeleteI know you've answered in pretty good depth, the question of how ATP is used in these reactions, but I was hoping to get an even deeper intuition for the answer.
From my current understanding, lets say that a reactant, R1, gets attached to a phosphate, and reacts with another reactant, R2, to make R2-R1-P
R1-Pi + R2 --> R2-R1-Pi
It appears to me that this, in itself, doesn't *really* explain why having that phosphate makes the reaction more favorable. But, if you are to posit that there's another step in the reaction, in which the product breaks as follows:
R1-Pi + R2 --> R2-R1-Pi --> R2-R1 + Pi
Now it would make sense that having Pi attached to R1 made it more reactive: because the product in the first step quickly broke down into R2-R1 and Pi, in accordance with Le Chatalier's Principle, the reaction gets driven to the right, thereby being 'favored' and producing more product.
Now, I'm not sure if this is how it actually works, but do you have any insights about my way of making sense of this?
Sorry for a very long comment.
VERY good post; easily described in a few short paragraphs what I couldn't piece together after reading fifteen pages of my text. Thank you.
ReplyDeleteI give below my logic of ATP driving an endergonic reaction. It is explained in textbooks as negative Gibbs energy of the overall reaction. This explanation has a problem. Gibbs energy is a function of state. Hence, the overall reaction if taken as sum of the two reactions, the endergonic, and the hydrolysis of ATP would not make sense, as it implies that the endergonic reaction does take place.
ReplyDeleteAt a given P,T there is an equilibrium ratio of number of product molecules to reactant molecules. You may like to look at it as product being continuously formed and breaking down to reactants, both being statistically a function of the number of collisions between molecules. If you add more reactants in an exergonic reaction, the number of product molecules increase and the reverse rate increases, till finally a ratio is reached where both rates balance out.
At equilibrium in physiological conditions, there are far more ADP than ATP. Hence when ATP is added, it gets hydrolysed to ADP and Pi. In an endergonic reaction, the products are fewer than reactants. Add ATP, and we get ADP and a new product, the phosphrylated reactant. This reaction is exergonic, but not as exergonic as ATP hydrolysis. The breakdown of the phosphorylated reactant to Pi and final product is also exergonic. Thus we have an overall greater ratio of final product to reactant. We do pay a price in terms of greater concentration of ATP to ADP, at an equilibrium when there is an endergonic product in the mixture.