Oxidative Stress Is An Electrochemical Concept That Focuses On An Inappropriate Rate Of Electron Transfer.

How damaging can electron transfer be if it is not regulated properly?

The electron is a particle that is responsible for transporting electric currents. Whenever we flick a light switch on at home, we’re connecting up a circuit that allows for the flow of electrons to power up a device that can then convert that electrical energy into light energy, such that our houses get illuminated.

In a chemical reaction, though, when two different chemicals come together, the transfer of electrons from Chemical A to Chemical B can result in a change in their physical and/or chemical properties.

That’s when the concepts of oxidation and reduction come in.

If Chemical A is able to extract electrons from Chemical B, we say that Chemical B is getting oxidised (loss of electrons), while Chemical A is a pro-oxidising agent that is getting reduced.

For example, steel is a type of metal that contains iron and can be used for multiple different applications. Some types of steel are prone to rust - and that occurs when the iron inside the steel comes into contact with atmospheric oxygen.

Oxygen, being an electron deficient chemical, readily extracts the excess electrons that iron metal has - in essence oxidising the iron metal into iron (III) oxide, which exists as a reddish-brown solid that is commonly defined as “rust”. Iron loses its electrons and becomes oxidised, while oxygen gains electrons and is reduced. Iron (III) oxide is a relatively stable compound that will not undergo further unnecessary oxidation or reduction reactions under ambient conditions.

But what if iron were to become chemically unstable upon oxidation, such that it then attempts to break up some other stable compound to reclaim whatever electrons that it has lost?

It turns into a highly unstable reactive oxygen species (ROS) that has to oxidise something else to extract electrons to maintain its stability. The propagation of these ROS oxidation reactions will generate even more unstable ROS species, which can also be termed as “free radicals”.

The concept of the redox potential

A stable form of Chemical B reduces the probability of ROS propagation. Different A-B pairings will result in the development of different reduction-oxidation (redox) potentials, as highlighted in the table below:

Table 1: A list of various redox pairs that can be found in the body

We can see that there are quite a fair few redox potential pairs in the table, all of which correspond to various biochemical nutrients that we can find in our diet or in our bodies.

The redox potential (E0) of any given redox pair indicates its oxidative capability. A more positive value of a pair indicates that it is more likely to oxidise something, while a more negative value of a pair indicates that it is more likely to be reduced.

If we were to look at Vitamin E oxidised/reduced at +0.37 V and Vitamin C oxidised/reduced at +0.08 V, what we would expect is that Vitamin C in its reduced form will get oxidised by Vitamin E in its oxidised form, such that Vitamin E will get reduced back into its reduced form. That’s the spontaneity of the reaction right there. The overall redox potential would be + 0.37 - (+ 0.08) = + 0.29 V, which is positive.

Reduced Vitamin C would be able to spontaneously reduce oxidised Vitamin E, and as such be able to regenerate any Vitamin E. Because Vitamin E exists as an antioxidant in the body, it will end up getting oxidised by any pro-oxidant ROS in the body.

Having a reduced version of an antioxidant exist as a lower redox pair in the body will help to regenerate any oxidised species higher up on the table.

Glutathione is considered to be the “master antioxidant” in the body because the cells in our body are capable of producing it and regenerating it (via the activity of the glutathione reductase enzyme). And we can see based on its position in the table that it can regenerate spent dietary antioxidants such as Vitamin C and/or Vitamin E all too easily.

Which brings us back to the regulation of electron transfer then.

What is the problem with a dysregulated electron transfer mechanism in our body?