Let’s have a refresher course on the basics of how each element deals with its own electrons and what influences the way electrons are dealt with in forming chemical compounds before we tackle how all the different lobes (representing electron probable locales) within a single energy level’s sub-shells can get close enough to allow their respective electrons to interact.

For this reason, we’ll use the more simple Bohr model to visualize the chemical compounds for salt and oxygen. In the salt compound, you will see that the chlorine atom is not very kind as it yanks away the one electron in the outermost (valence) shell in the sodium atom. So, there is a "hostile" transfer of an electron from the sodium to the chlorine atom.

In oxygen, one oxygen atom actually shares one electron from its valence shell electrons with the other oxygen atom, and vice versa. Now, why does a rude electron transfer occur in one scenario, while in the other, the atoms nicely share the electrons? The answer lies in two very important points: the importance of valence shells, and each atom’s number of protons and electrons.

We've learned that the importance of each electron's location within the atom depends on many factors. Yet its exact location is key to understanding how elements come together to form the compounds that sustain life.

And so it comes down to sub-shells. Why do they exist? What dictates their shape? And what else differs about them? As you can see in the image, each electron’s situation within an atom is a culmination of numerous factors, and explains why only certain electrons within the same energy level interact with electrons of other elements, whereas others do not.

Look at this chart depicting the various sub-shell shapes of the magnetic quantum numbers for s, p, d and f. What you are seeing are 3-dimensional regions of space representing the probable locations in which each electron for that atom can reside. Each sphere, teardrop, donut, or lobe is a region in which at any point in time you can locate one or two electrons. And so as an atom possesses more and more electrons, more regions need to be demarcated/allocated for each electron.

One mustn’t forget about all the other electrons and their residential 3D spaces! So the chart shows these spaces separated by their n and l quantum numbers. The upper right image shows how merely 3 of the sub-shells, 1s, 2s, and 2p, occupy/overlap within the same atomic space. It's pretty complex already, isn't it?

Next, we’ll see how the sub-shell of one element interacts with a sub-shell of another element!

WHY are we spending SO much time on electrons? You may be secretly begging me to cease, but I cannot, because electrons are the essential reactive particle of the atom. Whenever a compound is formed, like salt (NaCl), the two elements, Na and Cl, are bound together because of electrons! Thank you chemistry for making my food taste delicious.

Where the electrons are within the spatial dimension of the atom, and how much energy they do or do not possess, are at the heart of all chemical reactions. This is why we are making such an effort to figure out where the electrons are and how much energy they possess. These facts are key to understanding why some elements react more readily with others, why different kinds of bonds form between elements, and, why some chemicals undergo changes with additions of energy (in the form of heat), or with other chemicals. It’s all about the electrons.

So, to look at Aufbau, Hund and Pauli more closely, I’ve isolated one energy level in the previous graphic for us to scrutinize. Next, we’ll go onto looking at some mind boggling sub-shells.

I know these last few posts have started to go beyond the level of comfort, at least for me, but I invite you to bear with me and continue along this electron adventure. Because it's about to get worse!

I have presented the complex scenario that is the inner life of an atom. Furthermore, I have focused on how complex it is to BE an electron. Now, we finally get to the heart of the matter!

The calculation of an electron’s possible whereabouts boils down to 4 quantum numbers that are assigned to each electron, whereby no electron can have the same four numbers! What exactly are these numbers and what do they mean?

Quantum number 1 is the energy level of the electron and is designated as the “n” number. The higher the number, the higher the energy that electron possesses.
Quantum number 2 is the orbital angular momentum and is designated by the “l” number. In visual terms, it describes the shape of the cloud in which the electron is generally found.
Quantum number 3 is the magnetic quantum number, designated as “ml”. In visual terms it refers to the different sub-shells within an energy shell (hint, has to do with valence electrons. We will get into that in a future post).
Quantum number 4 is the spin magnetic quantum number “ms”. This describes the electron’s spin around its own axis, so as it’s moving about, it’s also spinning, like the Earth does, around its own axis.

We’re going to take this graphic APART and ADD to it. But you may be wondering, WHY are we spending SO much time on electrons? Can anyone answer that? Because it’s frankly the crux of chemistry, and life as we know it!I

One of the most important distinctions that moved the atomic model beyond Bohr's atomic model was defining exactly what an electron was. At first, it was assumed that if an electron carried a charge, then it must have mass and therefore must be a particle. And this is true. But other experimental data collected confirmed that the electron also exhibited behaviors of a wave, like light!

So, how does this impact upon describing the location and path of an electron? If we now know that an electron has an oscillating path defined by a wave-function, we begin to appreciate how complex the movement of each electron can get as it 1) runs in its own oscillating path, 2) feels the repulsive force from other electrons, 3) feels the attractive force from the protons.

Determining the orbits and locations of such minute particles reveals that locating an electron at any point in time is an impossibility. Erwin Schrodinger, after Louis deBroglie’s significant discovery of the electron’s dual nature as particle and wave, further defined the “orbital shells” as "probability-of-location clouds." Each shell is defined by energy level and further divided into “sub-shells,” shown in the graphic as 2s and 2p in Schrodinger's model. But it was becoming quite clear that a one-dimensional concentric planetary tracks (as in Bohr's model) were not sufficient to describe an electron’s position/path of the dual nature of an electron.

We've been getting into the nuts and bolts of what happens within an atom to an electron. Let’s do a theoretical exercise and start filling electrons into their orbital shells, from the lowest energy on up. The first 2 electrons can occupy the 1st orbital “shell,” but because they repel each other with equal force, they find the spot in the orbital that is equidistant from each other within the shell. If we stopped here, we can calculate the electron’s energy with quantum mechanics as Bohr proved. But with a more complex system like the Nickel atom, we can’t stop there.

Up to 8 electrons can occupy the next orbital shell. Though one electron wants to remain as far away from the next, there’s another electron in the way that is also dealing with repelling and being repelled. That is, instead of 1 electron exerting on and experiencing force from 1 electron as in the 1st shell, this 2nd shell electron will be exerting a repelling force, as well as having force exerted upon it, by 7 other electrons.

What's more, the shells aren’t isolated from one another! The 1st shell electrons exert force on the 2nd shell electrons etc. AND likewise those 2nd shell electrons are repelling the 1st shell electrons. So at this point, all 10 electrons are exerting and experiencing repelling forces from numerous electrons in their own orbital as well as from other occupied orbitals. In the lower left figure, I imagine visually what this might look like, using dotted lines of equal length to represent the equal distances represented by the equal forces of repulsion that each electron is experiencing and exerting. So in my map, I theorize how all 10 electrons might orient themselves to be most optimal with respect to another.

NOW, let’s fill in orbital shell 3, which can hold up to 18 electrons. In the right-hand figure, I try the mapping thing again, spreading these electrons out equally within their orbit, but can only fit 16 of the 18 remaining electrons into the orbital. If I simply increase the radius of the orbit, I can get all 18 electrons in the orbit! Et voila!

But no, this does not happen in reality. Why?

Before I move on, I want to make a clarification on valence shell capacity: 18 is the maximum occupancy of electrons that the 3rd shell can hold, but if that 3rd shell is the valence, or outermost shell, one is taught in school that the max. number it can hold is 8. While orbital "shells" are a concept that is no longer valid, for now, think of these shells as having "sub-shells," with the outermost sub-shell being the valence shell. Thus they determine that atom's chemical properties since they are the ones that interact with the valence electrons of another atom/s. Basically, it gets so much more complicated from here, and what I have been trying to do the last few posts is explain some reasons for why Bohr's atomic model falls apart. Check out: http://chemistry.tutorvista.com/inorga…/electron-shells.html