Electron Clouds
The Bioenergetic Blueprint #06
NB: This article is part of the Bioenergetic Blueprint series. The series is meant to be read in order. You can find its table of contents right here.
Friends,
Merry Christmas!
I hope you and your loved ones are enjoying yourselves these days. And what’s a better way to do so than to learn a little more about electron configuration? After all, at the core of what we want to understand in this series is how atoms interact with one another, especially in the context of our physiology and energy metabolism.
In the previous article, we mentioned one way to think about electron distribution is to do so in terms of “shells” or “layers”. Remember: electrons “want” to be closer to the nucleus since that’s where they’re most electronically stable, but as more of them join the party, they need to find ways to accommodate for each other. Thus, some of them end up closer to the nucleus (relatively lower energy), and some farther (relatively higher energy). This is often represented using Bohr’s model of the atom.
But there is another—more accurate—way to capture the true nature of electron behavior: not representing them through orbits, but orbitals—which, despite sounding similar, are very different things.
Whilst orbits suggest there is a fixed path electrons follow, orbitals understand they actually exist within somewhat-defined regions of space. According to quantum mechanics, it’s impossible to pinpoint both the position and momentum of an electron at the same time (see Heisenberg’s Uncertainty Principle), so we have to work with probabilities of finding electrons in certain regions of space around the nucleus.
Hence, orbitals are 3D wave functions where electrons are likely to be found at any given moment. More commonly, an orbital is said to be the 3D region of space neighboring the nucleus in which you can find an electron >90% of the time. For this reason, orbitals are also referred to as “electron clouds” or “probability distributions.”
Here’s what’s happening:
If you took a billion pictures of a helium atom (with its 2 protons, 2 neutrons, and 2 electrons), you’d notice that the nucleus would remain fairly still, but that its electrons would be in a different spot around the nucleus in each picture. And if you were to overlay all the pictures, you’d see that most of the time, the electrons would be found in a specific region or “cloud” around the nucleus.
That’s exactly what’s going on in the image above: we can see “superimposed pictures” of a helium atom across time. The black gradient represents the electron cloud where the electrons are most often found. The darker the color of a particular point, the more probable an electron is to be found there at any given moment.
Mind you: Orbitals exist only because we define them. There is no physical barrier constraining the electrons to that particular area, and they will sometimes be outside of it—but over time, that fuzzy region is where you’ll be most likely to find them.
Orbitals & Energy Levels
Orbitals are not always a simple spherical cloud. They can also have some pretty mind-bending shapes. This is not surprising considering electrons themselves are pretty mind-bending since their behavior follows what’s known as wave-particle duality. In short: electrons exhibit the properties of both particles in space with a specific mass and waves. This topic is the cornerstone of quantum mechanics.
Which orbital shape electrons form depends on two main variables:
The electron’s energy level which can be reframed as its potential gravitational energy vis-à-vis its atom’s nucleus. It is usually indicated by a principal quantum number (1, 2, 3, …), as we saw in the previous article.
The electron’s angular momentum (a measure of its rotation). It is usually indicated by an azimuthal or orbital quantum value (commonly s, p, d, f, g, h, …).
In simpler terms:
The number captures how big and far away from the nucleus the orbital gets.
The letter captures what type of shape the orbital has. The s orbital is spherical, the p orbital has a dumbbell shape, the d more of a cloverleaf shape, and after f things get pretty wild.
Here’s a table depicting the gist of this:
In the image, rows represent principal quantum numbers (1 to 7) and columns represent orbital types (s, p, d, f). Principal quantum numbers increase as you move down, indicating larger orbitals; and orbital shapes get weirder as you move to the right, indicating more complex angular momentum characteristics. Each individual cell in the grid shows a specific energy level at a specific orbital shape.
Filling Orbitals
Each orbital can house up to 2 electrons. Visually, we will represent one of these electrons with a half-arrow pointing upwards “↿”, and the other one with a half-arrow pointing downwards, “⇂”. These arrows are supposed to represent the electron’s complementary “spin”, a quantum mechanics property that explains why only 2 of them can occupy orbitals.1 When we see both (i.e., ↿ ⇂), we know there’s a pair of electrons fully filling up that orbital.
Now, when one orbital gets filled, additional electrons move to the next available orbital. This filling process follows a specific order based on energy levels, with electrons filling lower energy orbitals first, before moving to higher ones. In other words: the table is read from left to right, and top to bottom: 1s, 2s, 2p, 3s, 3p, 3d, etc.
For instance:
Helium has 2 electrons, so they can both reside in the 1s orbital.
Then, let’s consider lithium, which has 3 electrons. Because there’s no space for all 3 of them in the 1s orbital, the remaining electron will have to reside in the next available orbital: the 2s orbital.
Beryllium has 4 electrons, so it will fill out both the 1s and 2s orbitals with 2 electrons each
Boron, however, has 5 electrons, so the remaining electron will have to reside in the next available orbital, which is not the 3s orbital—because the electron is “not willing” to assume a higher energy level yet—but the 2p orbital.
Confusing? Just think of it this way:
Because electrons “don’t want” to exist at higher energy levels (here, 3), they first try to reside at the current energy level (here, 2) by squeezing themselves into weird shapes that will allow it. Once enough of them do this, then the next ones will be forced to go up to a higher energy level. It’s basically how when 6 people want to ride a 5-person car, they would rather be squeezed in the back than have 1 person go in the trunk.
Thus, non-s orbital have weird shapes, and because they are non-spherical, how they are oriented in space matters. So, whilst the table we shared above treats the 2p orbital as one single orbital, it has in fact three variations depending on the plane on which it is oriented, respectively called 2px, 2py, and 2pz.
Each of these hourglass-like figures is one single orbital. Think of it like a 3D version of a sinus wave function, with each of its sides being relatively positive/negative. When all three 2p orbitals exist together in a single atom, their electrons repel each other a bit, so they end up looking more like balloons, a bit like this:
Again: Each of those colored balloon pairs represents one single orbital which can house up to 2 electrons. It might seem like there are 6 orbitals in the image above, but there are only 3 (“double-balloon”) ones.
When it comes to filling p orbitals with electrons, they get distributed one at a time across each orbital before any receives a second electron. This means each p orbital (2px, 2py, 2pz) gets one electron first, and only after all are singly occupied does a second electron enter each orbital.
This method of electron filling, known as Hund’s Rule, minimizes electron repulsion and contributes to a more stable electron configuration within the atom. Filling it sequentially is crucial because incomplete orbitals—acting as sort of “electronic holes”—significantly affect an atom’s chemical reactivity and bonding behavior.
Let’s take the example of neon, for clarity.
Neon has 10 electrons. Therefore, it can house them in the following manner:
2 electrons occupy the 1s orbital. Easy.
2 electrons occupy the 2s orbital. Easy.
6 electrons occupy the 2p orbitals, filling them sequentially:
First, 1 electron each goes into the 2px, 2py, and then 2pz orbitals. That’s 3.
After each of these orbitals has 1 electron, they each receive a second electron. So, that’s 6.
This means all of the neon’s 1s, 2s, and 2p orbitals are fully filled. Its orbital diagram would therefore look something like this:
Neon is known to be a very stable, non-reactive chemical, which is why we categorize it as a noble gas. And from the logic above you can see why: All of neon's electron orbitals are fully occupied—with two electrons each—leaving no unpaired electrons or unfilled orbitals at its highest energy levels. Therefore, neon has no tendency to gain or lose electrons—hence its stoic or “noble” nature.
But what about carbon? Carbon has 6 electrons, so:
2 electrons occupy the 1s orbital.
2 electrons occupy the 2s orbital.
2 electrons occupy the 2p orbitals, filling them sequentially:
First, 1 electron each goes into the 2px, 2py, and—that’s it. We ran out of electrons.
This leaves carbon with just two half-filled 2p orbitals (and one of them missing entirely), so it looks a bit more like this:
Carbon’s half-filled 2p orbitals play a crucial role in its valence shell configuration, which is in turn crucial to its chemical behavior. The presence of unpaired electrons in its 2p orbitals makes carbon a perfect candidate to form covalent bonds with other atoms.
But how, exactly? If carbon only has 2 unpaired electrons in its 2p orbital, and the rest of the orbitals are full… how come we say—as we did in the previous article—that carbon has 4 valence electrons? In other words: how does carbon make 4 covalent bonds with other atoms, to fulfill the octet rule and achieve a more stable electron configuration?
To talk about this, we need to stop considering atoms in isolation, and we need to start looking at how they interact with one another across time. So, in the next article, we’ll introduce the concept of orbital hybridization, and we’ll begin to explore why it is important in the context of energy metabolism.
Until then, upwards, and merry Christmas once again.
Yago
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This is something I won’t get into here because, frankly, I’m not sure I understand it well enough… and because I don’t even know if it’s actually a thing.
Once again -- this is an OUTSTANDING effort !!
Although I still don't understand how possibly Carbon has four valence electrons ?!