Let me take you back in time. To a simpler time. To…high school chem. (Or maybe first-year of college chem? I don’t really remember.)
Atoms: they are the stuff things are made of. They consist of three important components:
- Protons: the positively charged particles in the nucleus (atomic core). Kinda like an atom’s DNA, they determine the chemical (elemental) identity (e.g., whether it’s carbon, oxygen, or something else).
- Neutrons: neutral (no charge) particles in the nucleus. If you got a core full of proton prima donnas, you need neutrons to help keep the peace and make sure the nucleus (or opera house) doesn’t disintegrate on itself via radioactive decay (or lies, deception and backstabbing). This becomes very important for the super heavy stuff, e.g., Uranium, which has 92 protons and even more neutrons. For now, we won’t really concern ourselves with these.
- Electrons: practically weightless, but equally as charged as the protons, albeit negatively. They are the reason why atoms stick together (or in some cases don’t). They surround the atomic core and are literally what make the bonds between us (if we were atoms).
Depending on the number of electrons and protons, the resulting atoms can either be neutral (equal numbers of both) or carry a net charge, which makes them ions. Electrons are obviously electrostatically attracted to the protons in the middle, but they surround the atom in specific groupings. Analogous to the Internet, an atom is not a truck onto which you dump a bunch of electrons. An atom is a series of orbitals.
This is also the reason why the periodic table is periodic.
Each orbital has a specific number of slots, and each energy level has a set of orbitals. The first orbital at any given energy level is the s-orbital, which has a maximum of 2 slots; the second (if applicable) is the p-orbital with a maximum of 6 slots. The electronic structure can be written out as level-orbital-occupancy (e.g. for Hydrogen, 1s1 = 1st level, s-orbital, just one spot). In the first level (H(ydrogen) and He(lium)), there is only an s-orbital, but at higher energy levels you got these p-orbitals, too, so you have to fill up the 2s and 2p spots (8), etc.
If you look at the first 18 elements of the periodic table, you see that for the electronic structure of an atom (not ion), you only stick in as many electrons as you need to match the atomic number (number of protons) so that the charge is neutral. Once you’re finish occupying a given energy level, you move to the next row or period of the table and do the same thing again. It kinda looks like this, but usually neater:
The atoms in each column or group have similar chemical properties, because their valence electrons, i.e., those at the highest energy level (at least at 0 K), determine their chemical behavior (how they bond). If you noticed, He could have been set right next to H, to sit above Be(ryllium) and M(a)g(nesium), which also have filled s-orbitals, but in the official table it got placed with the noble gases Ne(on) and Ar(gon), because it filled all the electron spots for its highest energy level (n=1). And that’s why it gets a crown, too.
At each subsequent level, the electronic structure can be written out with respect to that of the previous level’s noble gas. For example, anything in the second period (n=2), beginning with Li(thium), starts out as [He]2sx…, since it is understood that everything at n=1 is filled. Chemists have enough to worry about, so they decided to make things easier this way. I mean, when you do eventually get to something like Uranium, do you really want to write out the whole thing?
In summary, let us recall the classic Bohr model of the atom, where the nucleus is surrounded by rings (energy levels) of electrons: it is those electrons in the outermost ring that give rise to all the crazy things this stuff makes.
If only atoms were really so simple.
Want more details?
See: Clark, Jim. “Atomic Orbitals.” chemguide. n.p. Aug 2012. Accessed on 01 Jul 2015.
Want more drama, romance, and narcissism? Next time, in Act I!