Redox Reactions and Bioenergetics: How Your Body Knows Which Way a Reaction Goes

We have talked about delta G, the change in free energy, and why a negative delta G is what makes a biochemical reaction favorable. We related it to the equilibrium constant, to the reaction quotient, and to the idea that your body is constantly trying to move reactions forward in a way that extracts usable energy. But there is another way to look at this same question, and it is particularly relevant in metabolism because so many of the reactions your body runs are redox reactions.

Redox is short for reduction-oxidation. The “red” comes from reduction, the “ox” from oxidation. You could argue it should be called ox-red, but here we are. The important thing is that these reactions involve the transfer of electrons from one molecule to another, and that electron transfer is exactly how your body oxidizes glucose. You are quite literally rolling electrons down a hill, step by step, extracting energy at each step. Every major catabolic pathway, glycolysis, the TCA cycle, the electron transport chain, is built around this principle.

So the question becomes: how do you know which direction an electron is going to flow? How do you know if a given redox reaction is favorable? That is where redox potentials come in.

The Standard Hydrogen Electrode: Where the Numbers Come From

If you look at a table of standard half-cell reduction potentials, you will see values ranging from around positive 800 millivolts down to around negative 600 millivolts or lower. Every biochemical half-reaction has one of these values assigned to it. But where do these numbers come from, and what do they actually mean?

Everything is referenced to the standard hydrogen electrode, or SHE. This is the universal reference point in electrochemistry. The half-reaction it is based on is simple: two hydrogen ions plus two electrons yields H2 gas.

To understand why this is the reference, you need the Nernst equation. The Nernst equation describes the actual voltage of a half-cell reaction based on the concentrations of the species involved:

E = E° minus (RT / nF) times the log of Q

Where E° is the standard redox potential of the reaction, R is the gas constant, T is temperature, n is the number of electrons transferred, F is Faraday’s constant, and Q is the reaction quotient, expressed as the concentration of the reducing species over the oxidizing species.

For the standard hydrogen electrode under standard conditions, everything simplifies cleanly. The partial pressure of H2 gas is 1. The concentration of hydrogen ions is 1 molar. So Q is 1 over 1, which is 1. The log of 1 is 0. Multiply that by anything and you get 0. So the entire correction term drops out, and E equals E°. By definition, that value is set to exactly 0 millivolts. That is what the triple-equals sign means in this context. It is not measured. It is defined.

All other redox potentials in the table are measured relative to this reference. A more positive value means the reaction has a greater tendency to accept electrons (to be reduced). A more negative value means it has a greater tendency to donate electrons (to be oxidized). Once you understand that, the table starts telling you a story about the direction of electron flow throughout your metabolism.

What Changes When You Shift to Physiological Conditions

Standard conditions assume a hydrogen ion concentration of 1 molar, which corresponds to a pH of 0. That is obviously not what is happening inside a cell. Physiological conditions are closer to pH 7, meaning the hydrogen ion concentration is 1 times 10 to the negative 7 molar. That is a dramatically different concentration, and it changes the redox potential significantly.

Let’s run through the math. Plugging pH 7 conditions into the Nernst equation for the hydrogen half-reaction, the correction term becomes 59.16 millivolts divided by 2, multiplied by the log of 1 divided by (10 to the negative 7) squared. That works out to 59.16 divided by 2 times 14, which gives you approximately 414 millivolts. The correction term is negative, so the new redox potential under biochemical conditions is negative 414 millivolts.

If you look at a biochemistry table, you will see this value listed for the hydrogen half-reaction under what are called prime conditions, meaning adjusted to pH 7. The notation is E prime naught, and the value is negative 414 millivolts. The fact that this number shifts that dramatically just by changing the pH from 0 to 7 tells you something important: the concentrations of the species involved, particularly hydrogen ions, have a substantial effect on redox potentials and therefore on the thermodynamics of the reactions.

Connecting Redox Potential to Free Energy

The reason we care about redox potential is that it connects directly to delta G through a simple but powerful relationship:

delta G = negative n times F times delta E

Where n is the number of electrons transferred and F is Faraday’s constant (96.5 kilojoules per volt per mole). This equation tells you that if the change in redox potential is positive, delta G will be negative, and the reaction is thermodynamically favorable. If delta E is negative, the reaction does not proceed spontaneously in that direction.

This is not a separate principle from what we have already covered. It is the same thermodynamics expressed in the language of electrochemistry. A positive delta E means electrons are flowing spontaneously from a higher-energy donor to a lower-energy acceptor, which is exactly the situation in the electron transport chain. Each step releases energy because each electron carrier has a more positive redox potential than the one before it.

A Worked Example: Pyruvate to Lactate

Let’s make this concrete with one of the most important reactions in exercise physiology: the conversion of pyruvate to lactate by lactate dehydrogenase.

This reaction matters enormously during high-intensity exercise. When you are running glycolysis at full speed to produce ATP quickly, you accumulate NADH. The problem is that glycolysis requires NAD+ to keep running. If you cannot regenerate NAD+, the whole pathway stalls. The pyruvate-to-lactate conversion solves that problem by oxidizing NADH back to NAD+, allowing glycolysis to continue.

To understand why this reaction proceeds in the direction it does, look at the two half-reactions involved.

The first is the NAD+/NADH half-reaction: NAD+ plus 2 electrons plus a hydrogen ion yields NADH. This has a standard redox potential of about negative 320 millivolts.

The second is the pyruvate/lactate half-reaction: pyruvate plus 2 electrons plus a hydrogen ion yields lactate. This has a standard redox potential of about negative 185 millivolts.

Now combine them. To convert pyruvate to lactate, you need to reduce pyruvate. That means pyruvate is the electron acceptor, and NADH is the electron donor. To find delta E for the overall reaction, you take the reduction potential of the acceptor minus the reduction potential of the donor:

delta E = negative 185 minus negative 320 = positive 135 millivolts

A positive delta E means the reaction is favorable. Plug that into the free energy equation:

delta G = negative 2 times 96.5 times 0.135 = approximately negative 26 kilojoules per mole

Negative delta G, favorable reaction. NADH reduces pyruvate to lactate, and in doing so gets oxidized back to NAD+. The glycolytic machinery can keep running.

It is also worth noting that this conversion consumes a hydrogen ion. That is not just a bookkeeping detail. It means the pyruvate-to-lactate reaction has a buffering effect, consuming some of the protons that accumulate during intense exercise. This is one of the reasons the old narrative about lactate causing acidosis is incomplete. Lactate production is associated with acidosis, but it is not itself the cause.

Enzyme Kinetics: Why Thermodynamics Is Only Part of the Story

Thermodynamics tells you whether a reaction can proceed spontaneously. It does not tell you how fast it will proceed. That is the domain of enzyme kinetics, and for biological systems, the enzymes matter enormously.

Take the pyruvate-to-lactate reaction. If you dropped pyruvate and NADH into a beaker of water, nothing meaningful would happen. The reaction is thermodynamically favorable, but without lactate dehydrogenase there to lower the activation energy, the reaction proceeds too slowly to be useful. Enzymes are not changing the thermodynamics. They are making it practical for the thermodynamically favorable reaction to actually occur on a biologically relevant timescale.

Enzymes are also sensitive to the conditions they operate in, and this matters in the context of exercise.

Every enzyme has an optimal temperature at which it works best. Too cold and the reaction rate drops because molecules do not have enough kinetic energy to collide and react effectively. Too hot and the enzyme begins to denature, losing its shape and with it its function. During intense exercise, body temperature rises, and if it rises too much, enzyme activity across multiple metabolic pathways can be compromised.

Every enzyme also has an optimal pH. Hexokinase, which catalyzes the first step of glycolysis, operates near pH 7. The digestive enzyme chymotrypsin works at a much more acidic pH around 1.5 to 2. The reason pH matters is that hydrogen ions directly interact with the charged residues on an enzyme’s active site. Change the pH and you change the charge distribution, which changes the shape of the active site and how well it can bind its substrate. During intense exercise, as you hydrolyze ATP and accumulate metabolic byproducts, intracellular pH drops. That acidic shift alters the function of key metabolic enzymes and contributes to the decline in performance you feel as fatigue sets in.

Allosteric Regulation: Beyond Thermodynamics and Kinetics

There is a third layer of control beyond thermodynamics and basic enzyme kinetics, and it is what makes metabolic regulation truly sophisticated: allosteric regulation.

Allosteric regulators are molecules that bind to an enzyme at a site other than the active site and change its activity. They can either activate or inhibit the enzyme depending on the situation. AMP, ADP, and ATP are some of the most important allosteric regulators in energy metabolism. When AMP levels rise, indicating that the cell is running low on energy, AMP activates AMPK and also directly activates phosphofructokinase, the committed step in glycolysis, ramping up the pathway to produce more ATP. When ATP is abundant, it acts as an allosteric inhibitor of the same enzymes, signaling that energy is plentiful and the pathway can slow down.

This is worth distinguishing from the thermodynamic effect of high product concentrations. If you have a lot of ATP, Le Chatelier’s principle would predict that the equilibrium would shift back toward reactants, slowing the forward reaction. And that does happen. But the bigger and faster mechanism of inhibition is allosteric. When ATP binds to the allosteric site on phosphofructokinase, it changes the enzyme’s conformation and shuts it down much more directly and powerfully than a thermodynamic shift in the reaction quotient would alone.

The same principle applies to feedback inhibition more broadly. Products of a pathway can reach back upstream and inhibit the enzymes that produced them, preventing overaccumulation. Or, going the other direction, upstream signals can activate downstream enzymes before the substrate has even arrived, which is called feed-forward activation. Your body is not waiting for problems to develop and then reacting to them. It is anticipating demand and adjusting enzyme activity proactively.

Lactate: A Reversible Reaction, Not a Dead End

One final point about lactate that is worth making explicitly. The pyruvate-to-lactate reaction catalyzed by lactate dehydrogenase is reversible and highly dynamic. The direction it runs depends on the concentrations of pyruvate and lactate in the cell at any given moment, not just on the standard redox potential.

During intense exercise, when you are producing pyruvate faster than the TCA cycle can consume it and you need to regenerate NAD+ as quickly as possible, the reaction runs strongly toward lactate. But as exercise intensity drops or stops, as pyruvate is cleared into acetyl-CoA and shuttled into the TCA cycle, lactate can just as easily be converted back to pyruvate. The enzyme works the same way in both directions. It is the concentrations that determine which way it flows.

This is also why lactate is not waste product to be cleared as fast as possible. It is a fuel. The brain uses it. The heart uses it. And during exercise, muscles that are producing lactate are effectively passing substrate to other tissues that can oxidize it through the mitochondria. That concept, called the lactate shuttle, changes how you think about what is happening during high-intensity exercise and recovery.

Why This All Matters

The redox potential framework gives you a way to predict and understand the direction of electron flow throughout metabolism. It connects directly to free energy changes, which tells you whether reactions are thermodynamically favorable. And when you layer enzyme kinetics and allosteric regulation on top of that, you have a complete picture of how your body controls energy metabolism in real time.

Thermodynamics sets the rules. Enzymes determine the rate. Allosteric regulators tune the system to meet demand. All three have to be understood together to make sense of what is actually happening when you exercise, fatigue, recover, and adapt.

Next, we move into glycolysis, where these principles play out in detail across a ten-step pathway that your cells have been running for billions of years.