In the last lesson, we discussed the rules that valence electrons follow to dictate their bonding. But not all bonding is equal. There are some cases when the rules are broken.
Electrons will, generally, be shared evenly when the two atoms are the same or similar in terms of their electrostatic nature, or that within an atom which defines how strongly its positive charges (protons) attract negative charges (electrons). But what happens when one atom attracts electrons more strongly than the other?
There was another feature regarding the effective nuclear charge that we didn’t discuss. We now know that oxygen has an effective charge of +6. But what about the atoms it binds with? Well, from oxygen’s perspective, who cares? This is what makes oxygen so bad; because it has a high effective charge and a high-energy valence shell that it desperately wants to lower, it takes whatever is nearby.
The electron density in this bond is more heavily concentrated on the oxygen.
Image via University of British Columbia
We call the tendency of an atom to attract electrons electronegativity. Oxygen is extremely electronegative. When it comes into contact with something that is not as electronegative, like hydrogen, it will still “share”, but the electron cloud, or the possible areas that the electron could be, will be more dense near oxygen. It’s not enough that oxygen can share electrons to get into the party, but it behaves as if it’s the only atom that matters.
But, for as bad as oxygen is, it’s only the second most electronegative atom. When you take a look at that period of the periodic table, you see that oxygen has one atom between it and the next noble gas - fluorine. Fluorine has one more proton than oxygen, meaning it has a +1 increase in effective charge, making its nucleus exhibit a +7 force on outside electrons, and increasing its electronegativity.
Despite the electronegativity dictating a lot of the proceedings, chemistry is still a dance - the other participant matters.
How Shall We Dance?
So, what are some of the kinds of atoms that can form bonds? Let’s take a look at some of the most traditional groups.
Fluorine and Fluorine
I bring attention to this dance because most electronegative atoms are more than happy to bond with themselves. Nitrogen, oxygen, fluorine, chlorine, bromine and iodine are all diatomic atoms with high electronegativity.
Fluorine is missing only one electron in its valence shell. It’s fairly simple for it to bond with another fluorine and share one.
Hydrogen and Oxygen
The dance that most people are familiar with. The Earth is the domain of oxygen and hydrogen. Their dance produces water, otherwise known as H2O.
Remember, oxygen is missing two electrons in completing its orbital, while hydrogen is missing one. Therefore, two hydrogen atoms, attach themselves to one oxygen for the stability of all atoms involved.
Sodium (Na) and Chlorine (Cl)
You are aware of this one too. Sodium and chlorine combine to form sodium chloride - salt.
Sodium is extremely far away from its next noble gas. So, like, hydrogen, it would prefer to give its one valence electron away to another atom to be more like neon (Ne). What better kind of atom to give it to than an electronegative atom like chlorine?
Calcium and Calcium…and Calcium…
These atoms want to give up two electrons to be more like argon (Ar), the previous noble gas. But what about when there’s a bunch of calcium in the same place and no electronegative atoms around? Well, does it really matter if there are enough calcium atoms around to mimic the next noble gas, krypton (Kr)?
This kind of dance is strange, admittedly, because it’s not between two people, but some random number of participants.
Even though all of these combinations are valid for being similar to the noble gases, whether the next one or the previous one, they are not all the same. More specifically, the way each of these bond is completely different.
The first combination, fluorine and fluorine, is an example where electronegativity between the two atoms is equal. As such, they share their electrons equally. This kind of bond is referred to as covalent (co - mutual, valent - valence). Hydrogen and oxygen are also covalently bonded, but their electronegativity levels are different. This creates a unique situation in which most of the electron’s charge is isolated on oxygen, the most electronegative atom.
What you see here is the result of that electronegativity. The “δ” (delta) symbol represents that a particular atom behaves as though it’s positive or negative, according to the distribution of its electrons. In this case, oxygen and hydrogen are creating dipoles, which is a separation of positive and negative charges. Oxygen holds the negative dipole and both hydrogens each have a positive dipole. The presence of these uneven charges creates a unique kind of covalent bond called a polar covalent bond. You might be getting an inkling in the back of your mind about the possible implications of this polarity, but hold that thought, we’ll get to that in an upcoming lesson.
Sodium and chloride, unlike the other two combinations, are not covalent. That is, they don’t share at all. In fact, instead of the dance looking like two experts sharing a stage, it looks like one skilled dancer dragging a much weaker dancer along for the ride. This dance is referred to as the ionic bond, a bond between atoms at completely opposite ends of the electronegativity spectrum. The analogy makes complete sense when considering the effective nuclear charges. Chlorine, with its effective charge of +7 and a nearly-complete valence shell, can easily take possession of sodium’s electron, which has an effective charge of +1 and a willingness to get rid of an electron. Ionic bonds are a step up from polar covalent, where, instead of there being mere poles, there is one atom that is completely negatively-charged and one atom that is completely positively-charged, held together simply because negative and positive charges are attracted to each other.
Image via Annenberg Learner
Now, there is one more kind of dance, the most interesting kind, in my opinion. Has anyone wondered about what happens when there’s a bunch of atoms that want to give up their electrons, like sodium (Na), in one place without atoms that want to take them? Well, think about it this way: although sodium’s easiest path to being like a noble gas is giving up one electron, it’s not its only path.
First, let’s explain the dance. In this situation, all of the sodium atoms form metallic bonds with each other, which suits its nature of giving up electrons. Metallic bonds are a lattice, or net, of many atoms held together by an attraction to delocalized, or free, electrons from sodium.
This is not to scale. Remember, it’s a three-dimensional figure and electrons are everywhere.
Image by Keisho Inoue via IGCSE Chemistry
The delocalized electrons, essentially, form a very mobile cloud over the lattice, blocking the effective nuclear charge around sodium. But that’s not all - the sodium atoms being among so many electrons pushes it to be more like its next noble gas - argon (Ar) instead of its last - neon (Ne). Basically, its valence shell is being completed by the abundant sea of electrons surrounding it. For each individual sodium atom, this ends up being a pretty good deal because, if the rapidly-moving cloud is not over them (unlikely, but possible), it ends up being like neon since it’s successfully given its electron up to that cloud, but, if the cloud is over them, then it’s still stable, or, at least rapidly shifting between stability and instability. Boy, does taking “partially stable with a chance of stability” over “completely unstable” sound familiar in today’s times.
All of what we discussed dealt with one thing: the beautiful dances of chemistry. These dances create molecules, which are formed by two or more atoms bonding. In the cases where there are at least two different atoms, these molecules are also called compounds.
This natural effort by atoms leads to one thing: having the specific number of electrons, through any means, to become as stable as a noble gas. But what’s next? Is that all there is to atoms?
In a way, yes. If atoms (or molecules) are at their lowest possible energy-state, there’s no reason for anything else to happen. But there are many situations in which that is not the end.
However, that, my friends, is for later. Right now, there is another question to consider: with the seemingly infinite variety of molecules in our universe, how do you distinguish certain molecules from another? How, for example, do you distinguish water and rubbing alcohol? That…will be answered. Next time.
Quick Question: Are our bonding examples above all compounds? Are they all molecules?
As always, thank you for reading. You might be noticing that the scale of our discussions is increasing. It will only get bigger from here; make sure you join our mailing list to keep up with all future content.