Basically all of my posts as of late have been building up to metabolism, and this post is no different. Today, I’ll be talking about something pretty important to biosynthesis and catabolism: metabolic flux.
(My “x” key is broken, so if you like, you can think of this as a post on metabolic fluxzxx…)
All right, as my sister would say, let’s do this ish!
When we learned about ATP, we learned that it’s pretty thermodynamically unfavorable in comparison to ADP and AMP, but that doesn’t stop our cells from keeping it around. In fact, in our cells, the concentration is somewhere between 5 and 10 mM, much higher than ADP and AMP. This is called ATP’s “steady-state range,” and it’s pretty far away from equilibrium. So, how do we keep our ATP up this high? By catabolism, of course!
Catabolic metabolism is one side of the metabolic coin, and it’s pretty important. As you might know, or might have guessed from its name (“catastrophe!”), catabolism deals with breaking apart large molecules (such as carbohydrates, fats and proteins) to make ATP. This process makes CO2 and urea, waste products that our body disposes of, and helps us keep our levels of ATP way above what they really should be.
The flip side of this, of course, is anabolic metabolism. Anabolism does the exact opposite: it takes small molecules and builds them into complex biomolecules using ATP as an energy source. Anabolic processes consume large amounts of ATP, but in exchange we get all sorts of things that we really need to survive—hormones, muscle mass, DNA, etc., etc.
Obviously, cells have to maintain this cycle with a good amount of precision; they have to make enough ATP to fuel anabolic processes, which means maintaining something called “energy homeostasis.” This is a term referring to the balance of energy in a cell, but remember, this isn’t really an equilibrium; whether energy in a cell is “high” or “low,” it will always be higher than it would be at equilibrium proper.
Energy charge tells us how much energy content a cell has at a point in time. It’s defined by the concentration of ATP and half of the concentration of ADP divided by the concentration of all three adenine nucleotides (AMP, ADP, and ATP). (Only half of ADP is accounted for because it has half the number of phosphoanhydride bonds of ATP.) A high number means the cell has a high concentration of energy, and can carry out anabolism. A low one means cells will concentrate on catabolism in order to build up energy charge.
Okay, but what constitutes high and low? Unsurprisingly, an energy charge of about 0.9 (90%) is pretty good as far as the cell is concerned. When energy is that high, the cell can easily carry out anabolic metabolism (is in an anabolic state). Let’s make some hormones! Let’s make some glycogen! Let’s do all of the things that help our body to grow!
However, when you use up about 20% of that energy, your cell has a low energy charge. It switches over to catabolism, concentrating its efforts on improving that low energy level. And when I say your cell thinks it needs energy, I’m not talking, “Oh, I guess we should make some ATP, maybe,” either. A 70% energy charge has your cell convinced that it is dying oh my glob I need ATP liek, yesterday.
[ahem] I’ve had coffee.
Now, let’s take a step back and relate this whole business of fluctuating ATP concentrations to a larger picture that’s really important to metabolism. Let’s say you have a reaction like the one that converts ATP to ADP and Pi. How can you tell whether it’s spontaneous or not, given certain conditions?
(I know this seems redundant, but I promise, it’ll become relevant!)
Well, the equilibrium constant for any given equilibrium reaction is Keq, which is the concentration of products at equilibrium divided by the concentration of reactants at equilibrium. A high Keq means that products predominate at equilibrium; we’d expect this of the equilibrium between ATP and ADP + Pi, for example.
However, if you look at the ratio of the concentrations of products and reactants at a point when the reaction isn’t in equilibrium, you can predict which way the reaction will move. This relationship is described by the reaction quotient, Q. If Q is higher than Keq, that means that the concentration of products is higher than it normally is at equilibrium, and the reaction will shift towards the reactants. The opposite is also true.
In this example, we can see pretty easily that Q for the conversion of ATP to ADP and Pi in a cell is smaller than its Keq. How do we know this? Well, we know that ADP and Pi predominate at equilibrium; since ATP is in high concentration in a cell in comparison to ADP and Pi (even when a cell has low energy charge!), we can only assume that, in a cell, this reaction very much wants to move toward products. Always. That means Q must be smaller than Keq.
Okay, keep that in the back of your head. Now we’re going to do a touch of thermodynamics.
Remember all of that talking I did about free energy change in my post about ATP? Well, now we’re going to look at that in more detail. If you’ve been introduced to the concept of Gibbs free energy already, you might remember this relationship:
∆G° = – RT ln Keq
That’s a jumbled up mess, so let me do some ‘splaining. First of all, ∆rG° refers to the free energy of a reaction at equilibrium (standard free energy change). This is going to tell you whether at equilibrium, you have more products or reactants. If Keq is a big number (products > reactants), the natural log of it will be a positive number, and the whole term will be negative. If the Keq is a small number (reactants > products), the natural log will be negative, and the whole term will be positive. This means that negative standard free energies favor products.
But what if our reaction isn’t at equilibrium? Well, in that case, you have to deal with a different equation that will take into account those nonstandard conditions. The equation for that looks like this:
∆G’ = ∆G° + RT ln Q
If we substitute for ∆G°, we get this:
∆G’ = – RT ln Keq + RT ln Q
Quite a chunk of an equation, eh? Let’s break it down a little bit. What this is telling us is, if your reaction quotient is bigger than your Keq, ln(Q/Keq ) (Remember your log identities? If not, don’t worry, I don’t know anyone who does, ehehe…) is a large number, and the whole term is positive. The opposite is true, too. That tells us that, if the ratio of products to reactions is bigger than it would be at equilibrium, the reaction will want to shift toward reactants, and ∆G’ will be positive. The opposite works, too—if you have fewer products than you would at equilibrium, the whole shebang will shift toward products, and your ∆G’ will be negative, to boot. ^^
So, essentially, you can conceptualize it this way: ∆G’ is a measure of how much energy would be released from a reaction that isn’t in equilibrium if it moved to equilibrium. The further away from equilibrium a reaction is (in either direction), the more energy it stores—the only difference is the direction the reaction has to move to reach equilibrium. The larger the difference between Q and Keq, the more energy stored in the reaction.
In other words, keeping reactions away from equilibrium is a way to store energy. That, my friends, is what generates what we call metabolic flux.
Now, let’s revisit that whole bit about cellular ATP concentrations. Like we said before, under standard conditions, ATP would be the least abundant of the three adenine nucleotides at equilibrium. Under those conditions, hydrolyzing a mole of ATP to product a mole of ADP and Pi would product -37 kJ/mol (∆G° = -37 kJ/mol). However, under cellular conditions, this conversion is way out of equilibrium; using the equation above, we find that the energy released when ATP becomes ADP and Pi under cellular conditions is -54 kJ/mol (∆G’ = -54 kJ/mol). That’s a 50% increase in the energy we get from ATP just by keeping it in abundance!
We can look at an example to see how this really works. Let’s say we want to make glycerol-3-Pi from glycerol and Pi using ATP. Now we have to adjust our calculations to account for the greater energy that we get from ATP. However, the energy of condensation for the reaction of glycerol and Pi under standard conditions is 9.2 kJ/mol, but since glycerol-3-Pi is kept in a fairly high concentration in cells, it takes more energy to make it in a cell (26 kJ/mol) than it would otherwise. Still, this reaction is spontaneous; the net free energy change is -28 kJ/mol under cellular conditions.
Metabolic flux is a pretty common way of driving metabolism. We can revisit our example of phosphocreatine, in fact, to see why this is a good way to store energy in muscle cells. Usually, muscular phosphocreatine levels are kept at 32 mM or so, while ATP is kept around 8 mM (and regular creatine is kept at 7 mM). Under these conditions, formation of ATP from phosphocreatine is favored, since the reaction is out of equilibrium with more reactants than usual. About 10 kJ/mol would be released for this conversion.
However, when there’s a lot more ATP in muscle cells than phosphocreatine, metabolic flux pushes the reaction in the other direction; formation of phosphocreatine from ATP becomes favorable, releasing 24 kJ/mol. Thus, muscles maintain their own little energy reserves by making clever use of metabolic flux.
As we can see, this keeping reactions out of equilibrium can be quite useful. However, it isn’t necessarily required for every step in a metabolism. For example, in the first stage of glycolysis, the reaction that starts the pathway is subject to metabolic flux (a phosphate is attached to glucose), but the next one (the isomerization of glucose-6-Pi to fructose-6-Pi) is at equilibrium. The entire second stage of glycolysis, too, consists of reactions at equilibrium, as does the non-oxidative portion of something called the pentose phosphate pathway.
We should also note that there are two factors that affect metabolic flux. First, to have metabolic flux, the reaction in question must be out of equilibrium. Second, there must be enzymes to mediate the reaction. This is because, even though it might be thermodynamically favored (the formation of products might be energetically favored), the reaction probably isn’t kinetically favored (the activation energy of the reaction is likely very high).
In fact, the modulation of enzymes is an important part of regulating different metabolic pathways. For example, for someone who doesn’t drink a lot of alcohol, the levels of dehydrogenase enzymes in their liver meant to deal with alcohol are kept relatively low. On the other hand, someone who drinks more alcohol regularly will have higher dehydrogenase levels. If these two people both drink a significant amount at one time, the person who usually doesn’t drink is at a higher risk for alcohol toxicity, since their liver enzymes are adjusted to a diet low in ethanol.
This would be the part where my professor would say, “Listen to me, kids: don’t screw with your liver.” However, the liver does even cooler things than this. Now that we’ve ended our discussion on the mechanisms of metabolism in general, we can get to talking about some of them—cool processes such as glycolysis and gluconeogenesis, for example.
This was shortish, right? …. Right?