exploring oxygen
I’ve been reading David Beerling’s fascinating but demanding book The Emerald Planet, essentially a history of plants, and their contribution to our current life-sustaining atmosphere, and it has inspired me to get a handle on atmospheric oxygen in general and the properties of this rather important diatomic molecule. Demanding because, as always, basic science doesn’t come naturally to me so I have to explain it to myself in great detail to really pin it down, and then I forget. For example, I don’t have any understanding of oxidation right now, though I’ve read about it, and probably written about it, and more or less understood it, many times. Things fall apart, and then we fall apart…
Okay, let me pull myself together. Oxygen is a highly reactive gas, combining with other elements readily in a number of ways. A bushfire is an example of oxidation, in which free oxygen is ‘consumed’ rapidly, reacting with carbon in the dry wood to produce carbon dioxide, among other gases. This is also called combustion. Rust is a slower form of oxidation, in which iron reacts with oxygen to form iron oxide. So I think that’s basically what oxidation is, the trapping of ‘free’ oxygen into other gases or compounds, think carbon monoxide, sulphur dioxide, hydrogen peroxide, etc etc. Not to mention its reaction with hydrogen to form water, that stuff that makes up more than half our bodily mass.
Well, I’m wrong. Oxidation doesn’t have to involve oxygen at all. Which I think is criminally confusing. Yes, fire and rust are examples of oxidation reactions, but so is a reaction between hydrogen and fluorine gas to produce hydrofluoric acid (it’s actually a redox reaction – hydrogen is being oxidised and fluorine is being reduced). According to this presumably reliable definition, ‘oxidation is the loss of electrons during a reaction by a molecule, atom or ion’. Reduction is the opposite. The reason it’s called oxidation is historical – oxygen, the gas that Priestley and Lavoisier famously argued over, was the first gas known to engage in this sort of behaviour. Basically, oxygen oxidises other elements, getting them to hand over their electrons – it’s an electron thief.
Oxygen has six valence electrons, so needs another two to feel ‘complete’. It’s diatomic in nature, existing around us as O2. I’m not sure how that works – if each individual atom wants two electrons, to make eight electrons in its outer shell for stability, why would it join with another oxygen to complete this outer shell, and then some? That makes for another four electrons. Are they now valence electrons? Apparently not, in this stable diatomic form. Here’s an expert’s attempt to explain this, from Quora
For oxygen to have a full outer shell it must have 8 electrons in it. But it only has 6 electrons in its valence shell. Each oxygen atom is actively seeking to get more electrons to complete its valence shell. If no other atoms except oxygen atoms are available, each oxygen atom will try to wrestle extra valence electrons from another oxygen atom. So if one oxygen atom merges with another, they “share” electrons, giving both a full outer shell and ultimately being virtually unreactive.
For a while this didn’t make sense to me, mathematically. Atomic oxygen has eight electrons around one nucleus. Six in the outer, ‘valence’ shell. Molecular oxygen has 16 electrons around two nuclei. What’s the configuration to make it stable? Presumably both nuclei still have 2 electrons configured in their first shells, that makes 12 electrons to make for a stable configuration, which doesn’t seem to work out. Did it have something to do with ‘sharing’? Are the shells configured now around both nuclei instead of separately around each nucleus? What was I missing here? Another expert on the same website writes this:
[The two oxygen atoms combine to] create dioxygen, a molecule (O2) in which both oxygen atoms have 8 valence electrons, so they are happy (the molecule is stable).
But what about the extra electrons? It seems I’d have to give up on understanding and take the experts’ word, and I hate that. However, the Khan academy has come to the rescue. In video 14 of his chemistry series, Khan explains that the two atoms share two pairs of electrons, so yes, sharing was the key. So each atom can ‘kind of pretend’, in Khan’s words, that they have eight valence electrons. And this is a covalent bond, unlike an ionic bond which combines metals with non-metals, such as sodium and chloride.
Anyway, moving on. One of the most important features of oxygen, as mentioned, is its role in water – which is about 89% oxygen by weight. But how do these two elements – diatomic molecules as we find them in our environment – actually come together to form such a very different substance?
Well. According to this video, when H2 and O2, and presumably other molecules, are formed
electrons lose energy to form the new orbitals, the energy gets away as a photon, and then the new orbitals are stuck that way, they can’t undo themselves until the missing energy comes back.
This set me on my heels when I heard it, I’d never heard anything like it before, possibly because photon stuff tends to belong to physics rather than chemistry, though photosynthesis rather undoes that argument…
So, sticking with this video (from Brigham Young University Physics Department), to create water from H2 and O2 you need to give them back some of that lost energy, in the form of ‘activation energy’, e.g by ‘striking a match’. The video turns out to be kind of funny/scary, and it again involves photons. After the explosion, water vapour was found condensing on the inside of the glass through which hydrogen was pumped and ignited…
Certainly the demonstration was memorable (and there are a few of these explosive vids online), but I think I need more theory. Hopefully I’ll get back to it, but it seems to have much to do with the strong covalent bonds that form, for example, molecular hydrogen. It requires a lot of energy to break them.
Once formed, water is very stable because the oxygen’s six valence electrons get two extras, one from each of the hydrogens, while the hydrogens get an extra electron each. The atoms are stuck together in a type of bonding called polar covalent. Oxygen is more electronegative than hydrogen, meaning it attracts electrons more strongly – the negative charge is polarised at the oxygen, giving that part of the molecule a partial negative charge, with a proportional positive charge at the hydrogens. I might explore the effects of this polarity in another post.
The percentage of oxygen in our atmosphere seems stable at 21% – that’s to say, it appears to be the same now as when I was born, but that’s not a lot of time, geologically. The issue of oxygen levels in our atmosphere over geological time is complex and contested, but the usual story is that something happened with the prokaryotic life forms that had evolved in the oceans billions of years ago, some kind of mutation which enabled a bacterial species to capture and harness solar energy. This green mutation, cyanobacteria, gave off gaseous oxygen as a waste product – a disaster for other life forms due to its highly reactive nature. The photosynthesising cyanobacteria, however, multiplied rapidly, oxygenising the ocean. Oxygen reacted with the ocean’s iron, creating layers of rust (iron oxide) on the ocean floor, later visible on land through tectonic forces over the eons. Gradually over time, other organisms evolved that were adapted to the new oxygen-rich atmosphere. It became an energy source, which in turn produced its own waste product, carbon dioxide. This created a near-perfect cycle, as cyanobacteria required CO2 as well as water and sunlight to produce oxygen (and sugar). Other photosynthesising water-based and land-based life forms, plants in particular, emerged. In fact, these life forms had harnessed cyanobacteria as chloroplasts, a process known as endosymbiosis.
I’ll end this bitsy post with the apparent fact, according to this Inverse article, that our oxygen levels are actually falling, and have been for near a million years, and that’s leaving aside the far greater effects due to human activity (fossil fuel burning consumes oxygen and releases CO2). Of course oxygen is very vastly more abundant in the atmosphere than CO2, and the change is minuscule on the overall scale of things (unlike the change we’re making to CO2 levels). It will make much more of a difference in the oceans however, where the lack of dissolved oxygen is creating dead zones. The article explains:
The primary contributor to these apocalyptic scenes is fertilizer runoff from agriculture, which causes algal blooms, providing a great feast for bacteria that consume oxygen. The abundance of these bacteria cause O2 levels to plummet, and if they go low enough, organisms that need it to survive swim away or die.
Just another of the threats to sea-life caused by humans.
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