What does the phrase "reducing atmosphere" mean in quantitative terms?

Question

There is a debate about "how reducing" the atmosphere of the early Earth was, and this question is my attempt to grasp toward an understanding of what that that phrase really means. As a bit of context, the Miller–Urey experiment produces a wide variety of complex organic molecules, but people say it requires a very reducing atmosphere to work, and it's disputed whether the early Earth's atmosphere was reducing enough for analogous processes to occur. Ultimately I would like to understand what it is about reducing atmospheres that allows complex organic molecules to be created so easily, but the first step is to understand what it really means to say one mixture of gases is "more reducing" than another.

If anything can be more reducing than something else, it seems like $\ce{CH4}$ gas should be more reducing than $\ce{O2}$, so I'm trying to understand the extent to which this can be quantified. However, I'm currently self-learning chemistry, and I'm having difficulty seeing connections between some of the concepts. (So apologies in advance for the rambling nature of this question.)

It seemed like one place to start might be to draw the net reaction of methane oxidation:

Obviously, if you add up the total oxidation state of every atom, you get zero - the carbon gets oxidised, the oxygen gets reduced, and the hydrogen stays the same. This must always be the case for any reaction (right?) since electrons are conserved. At first this made me think that maybe you can't really say that one thing is objectively "more reducing" than another. Though of course you can still talk about its tendency to oxidise or reduce carbon, and maybe that's what people are talking about when they say an atmosphere is "reducing" or "oxidising".

But then I came across the concept of redox potential, which does seem to be about things being objectively more reducing or oxidising than others. The problem is, the discussions of the concept that I can find online seem to be in terms of electrons "being transferred" from one molecule to another. But in the reaction above, as far as I can see, electrons aren't being transferred from the $\ce{O2}$ to the $\ce{CH4}$ but instead both molecules are being destroyed and replaced with completely different ones. So I'm having trouble seeing whether, or how, the concept of redox potential can be applied in this case.

Given this confusion, I guess my specific questions are:

1. When people talk about an atmosphere being "more reducing" than another, are they referring to redox potential or something else, such as the degree to which the atmosphere reduces or oxidises carbon? If it's the latter, is it something that can be quantified, or is it more of a subjective judgement?; and

2. What is the relationship between the concepts of the redox potential of a substance on the one hand, and the oxidation states of atoms on the other? Can the concept of redox potential be applied to gases like $\ce{CH4}$ and $\ce{O2}$?

Answer 1

I’m an environmental chemist and we frequently refer to ‘oxidizing’ or ‘reducing’ conditions in the environment. I’ll use soils as an example. The ‘oxic zone’ is the area of soil that oxygen can reach by diffusion and is seen as being under ‘oxidizing’ conditions because the reduction of $\ce{O2}$ is the dominant redox reaction that occurs in that zone. Elements that have multiple oxidation states tend be in the higher oxidation states. For example, sulfur will exist primarily as sulfate ($\ce{SO4^2-}$). As you get further away from the Earth’s surface, the oxygen is depleted (by microbes) and the redox conditions change: the dominant reduction reaction changes. Usually in this order: $\ce{O2}$, nitrate, manganese(III,IV), iron(III), sulfate and finally $\ce{H2}$. When the conditions become so reducing that the microbes are using $\ce{H2}$ as an electron acceptor, they making methane via a process called methanogenesis. This order arises because of reduction potentials. Microbes get the most energy for the least effort from reducing $\ce{O2}$, so they do that until it runs out, which is when they switch to the next best thing available.

Ultimately what it means to be a ‘reducing environment’ is that if you added a compound with a higher reduction potential then the dominant redox pair, it would become reduced. For example, if you could put $\ce{SO4^2-}$ in a methanogenic environment, it would be reduced to $\ce{H2S}$ very quickly.

I’ve read that the composition of Earth’s atmosphere in pre-life earth was very different from today. If I recall correctly, there was no $\ce{O2}$ as it was formed by microbes later on. Without $\ce{O2}$ in Earth’s atmosphere, it makes sense that conditions would be very reducing based on an analogy to the soil redox zones.

Answer 2

I'm not certain mine will be the best answer, but I may be able to get this started.

The short answer is yes, redox potentials are the key to determining how reducing or oxidizing a reaction is. What your simple balanced redox equation (which is correct) does not show are the various half reactions that produce the complete reaction.

The big oxidizing half reaction in your equation is the classic reduction of oxygen:

$\ce{2O2 +8H+ + 8e- ->2H2O}, E_0=+1.23V$

I've already balanced this half reaction for your combustion equation. The positive E0 value of 1.23 volts says this reaction tends to occur spontaneously (given some activation energy).

However, the other half reaction, the oxidation of carbon in methane, is likely MUCH more complex, because it involves many steps. Here's where my answer my not be completely correct. Below is one possible sequence of half reactions:

$\ce{CH4 +H2O -> CH3OH + 2H+ + 2e-}, E_0=-0.50V$ $\ce{CH3OH -> HCHO + 2H+ + 2e-}, E_0=-0.13V$ $\ce{HCHO +H2O -> HCOOH + 2H+ + 2e-}, E_0=+0.03V$ $\ce{HCOOH -> CO2 + 2H+ + 2e-}, E_0=+0.11V$

Notice first that the half reaction oxidizing methane to methanol has a negative potential of -0.50 volts and so is not spontaneous; it needs an oxidizer like oxygen to get started. In the sequence of half-reactions I list, the potentials get progressively bigger, so that by the time we are oxidizing formic acid to carbon dioxide, the half reaction is slightly spontaneous (positive potential).

So this shows that you were correct that the reaction involves the creation and destruction of various compounds, but the sequence (the mechanism) depends on redox reactions which DO transfer electrons.

Finally, the reason my answer may not be best is that the mechanism I propose may be wrong. While all these half reactions do exist and produce a spontaneous total, the total potential sums to 0.74 volts. At least one source I read suggested that the actual potential of methane oxidation to carbon dioxide is 0.92 volts. So there may be some other sequence of half reactions, which I could not track down, which is a better mechanism than the one I propose. Still, I think my answer gets the principles right.