Act I: Give and Take – Laws of Ionic Attraction

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*Sigh*

Noble gases — they just don’t care to bond with anyone.

00_argon

Since they filled up their valence band, they have no affinity to add more or get rid of any electrons, and are happy just to stay the way they are.

00_valenceband

Most other elements, however, feel very differently about their neutral electronic state.

01_wannabe

As a result, these atoms tend to bond over their lamentable situation. Ionic bonding just happens to be one way to do it. This requires the atom to ionize, or basically gain or lose electrons and carry a net charge. What an atom is likely to do basically comes down to the attraction and repulsion between charged particles.

But wait, you say, shouldn’t they all just be happy? If you have the same number of protons and electrons, then all should be well and everyone should be happy.

No, no, no, I respond, it’s all about distance! All protons are in the core, and all electrons are in their orbitals; the higher the energy level of the orbital, the further it is from the core and the more loosely bound its electrons are by the nucleus. Plus, the more electrons already attached to the atom, the harder it can be to stick another one in, because all the others are quite literally repulsive.

See, the noble gases are in a fantastic position: they are Goldilocks. Take Ne(on) and its neighboring elements as examples:

  • F(luoride) has a lot of protons to hold its electrons, and it could totally take on one more in its 2p-orbital. It will release energy as a result, which is known as the electron affinity energy. In contrast, if you add another electron into neon, it can’t take it in 2p (full) and would need to stick in 3s, which is not only further away from the nucleus but also at a higher energy level. Like trying to feed broccoli to a five-year-old, possibly after giving him/her pizza, you either need to exercise brute force or simply accept that it’s not going to happen.
  • On the other side is sodium (Na). Na happens to have a single lone electron in a far-away orbital, and that electron could easily disappear with a flick; this process can be done by putting in just a little energy, known as the ionization energy. This is the case even though Na has one more proton than Ne, because at least with Ne, the outermost electrons are much closer to the nucleus. Similarly, once Na becomes Na+, it’s nearly impossible for it to become Na2+, because it will already be a positively charged bad@$$ with closely held electrons, and you do not want to mess with that.

In summary, for most non-noble gas elements, ionization is pretty likely.

There are quite a few well-known ionic solids, including in the realm of bling: sapphire and ruby are colored versions of the mineral corundum (Al2O3); and spinel is the result of oxygen ionically bonding with magnesium and aluminum (MgAl2O4). In addition, salts are a well-known group of compounds resulting from ionic bonding; their distinguishing feature is their ready ability to dissolve in water. Some common examples include:

  • Sodium bisulfate (NaHSO4) and potassium dichromate (K2 CrO7), which are used in photographic bleach;
  • Magnesium sulfate (MgSO4), a.k.a. epsom salt;
  • Calcium carbonate (CaCO3), which occurs naturally as chalk, limestone, marble, etc.

Of course, the most famous (and arguably simplest) salt of all is….salt.

01_anticlimactic

So how does ionic bonding in salt occur? Well, when a sodium (Na) atom, who has one too many electrons….

01_sodium

…. meets a chlorine (Cl) atom, that has one too few…

01_chloride

(Seriously, have you ever tried hula-hooping with an open loop? It just doesn’t work.)

…they fall in love, talk about life together as a compound, and finally tie the ionic knot.

01_proposal

As a result, Na becomes the cation (Na+), and chlorine becomes the anion (Cl), and both of them get their noble gas-esque, Ne-or-Ar-like electronic structures. If this isn’t chemistry, I don’t know what is.

I suppose you could say they’re “joined at the hip”, but that seemed like an extreme analogy.

01_hysteria

So let’s say they’re three-legged racers instead.

01_threelegs

Besides, it’s a pretty good metaphor for the whims of a relationship, no? I mean, compromise and agreement on which foot to move forward is really key to progress. Although I guess if you’re persistent enough, you could get your way anyway…

01_party

Fortunately, because sodium chloride (NaCl) is an electrolyte (salt), can dissociate again in certain situations, such as when being dissolved in water.

01_swimming

However, they will continue to exist as ions in water and make it taste funny (or delicious, if you’re cooking!). In fact, the whole reason why water can even pull them apart is in part due to its own electronic polarity, but that’s a story for another day. Given enough time, water will eventually evaporate, leaving the salt crystals to themselves once again.

Please do note, that while you can throw as much water as you want on sodium chloride, and you can use chlorine to clean water, sodium is the last thing you want near water; it will explode, and that’s as good of a reason as any to keep it away.

In short, while the ionic bond is defined as a “give-and-take” relationship, it really isn’t so lopsided at all. In the end, everyone is happy! Just as well, as there are different forms of love, so are there different types of bonding. Stay tuned for Act II!

Sources (accessed on 01-02 Jul 2015):
Ionization Energy and Electron Affinity.Purdue University Division of Education. Purdue University.
Nave, Carl. “Common salts.Hyperphysics.
Some useful salts.Eduvee. Edyia Ltd.
Spinel.mindat.org. Hudson Institute of Mineralogy.

Prelude: Chemistry Repeats Itself

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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:

IMG_0742

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.

IMG_0765

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!

A brief intro to the structure of bling

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Oh, crystals! I do love my crystals.

They’re the foundation material for most of my works, but it was not really until I was an undergrad that I really appreciated them.

This was partly because I had not realized that so many of the materials I dealt with on a regular basis were crystalline.

The thing is, when you have a solid material, the atoms have to sit somewhere, but where do they go?! Well, it’s not always easy to answer this, but if you have a general idea of what the atoms are, what sort of traumatic treatments (extreme heat, natural disasters, etc.) they had to face, what other junk got tossed into the mixture, and what the ambient temperature and atmosphere are, you could make some pretty good guesses. For centuries, this is exactly what people did!

Given enough time, energy, and a certain affinity for one another, with all atoms being more or less equal, they will try to sit in an ordered manner with one another. Many of my favorite materials happen to sit in a nice cubic structure, the simplest case being a simple cubic configuration, where one atom occupies each corner:

SC unit cell
It doesn’t get simpler than this.

But that’s just 8 atoms, and what do you care about 8 atoms? That’s practically nothing (seriously)! Well, if you keep propagating that cubic pattern in all the main directions along the cube’s edges (forward-backward, left-right, up-down), you’re going to get a larger crystal.

Crystal building
Yes, this is exactly how crystal structures propagate.

How large you ask? Well, up until you run into a wall.

wall
Well, I guess we’re camping out here. Indefinitely.

And if your atoms just don’t have enough time, energy, or discipline, they just are a mess.

Here’s how this manifests in some real materials you might actually care about:

  1. Some solids are obviously crystalline: diamond, quartz, sapphire, etc. You can tell because they have nice flat surfaces or facets that naturally occur when you try to cut these things into smaller pieces.
  2. Some solids are not so obviously crystalline: metals are the big one. Most metals we deal with are just collections of very tiny crystallites packed tightly enough that you can’t tell the difference between one or the other unless you have a microscope. So unless you’re dealing with something really REALLY special, that hunk of metal you’re working is polycrystalline.
  3. Some solids are not crystalline at all: glass is not just the name of a material but also a structural term. It just so happens that glass is a glass. How about that?

For (1) and (2), you got more or less long-range order, it’s just that for (1) it’s definitely longer than for (2). (2) happens when you have two armies running up against each other to do battle at their boundaries, and yeah, it’s a mess.

As for (3), this is what is known as “amorphous”. I mean, it kind of does have order, it’s just short-range order (OMG, the name of the blog in a blog post!!). It could be just a mess of cubes or polyhedrons randomly oriented with respect to each other, so order is limited to the nearest atoms only (e.g. as far as the cube corners go).

Short-range order
PARTY!

Or it really could be a complete mess, like freezing in time the chaos that ensues when throwing a bunch of primary schoolers into a classroom and ordering them to sit in their seats while candy rains from the ceiling.

Even shorter-range order
CANDY PARTY!

Clearly, order takes time to restore. Likewise, cooling silicon dioxide (SiO2) too quickly from a liquid to a solid does not give time for the atoms to settle into place (like it would in quartz), so it remains chaos (i.e., glass). But frozen.

Given enough time, the sugar and natural hyperactivity will wear off, and all the atoms will settle into their lowest energy state.

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The end.