Hello, everyone! I’m back in a surprisingly timely fashion! Weird, right? I know! My motivation stems from the fact that I now have a plethora of sticky notes on the wall of my study space, so I now study in decisive chunks instead of haphazard stints. Today I’m bringing you a review of some Inorganic I material that we covered last week: characteristics of covalent materials. This was the beginning of what’s going to be a long study of periodicity, but I figured since the material covered here is a bit unlike all the other stuff, I’d classify it differently.
(Sorry, my study music is Japanese…)
Allotropes are compounds composed entirely of single elements. Some of them are substances that we are familiar with from day-to-day life; graphite and diamond, for example, are both allotropes of carbon, oxygen and ozone are allotropes of oxygen, and neon gas is an allotrope of neon. However, there are a great many that we probably haven’t encountered that will still give us great insight into the properties of elements in different columns of the periodic table.
So, how do we start our study? Why, systematically, of course! We’re going to start to the far right of the periodic table and work our way left until we run into metal!
(All of what we’ll look at here concerns nonmetals with an oxidation state of zero, so there’s no electronic funny business afoot.)
Group 18/8A—The Noble Gases
The noble gases are found in the very last column (group) of the periodic table, and they have full valence shells. These are so called because, while they can be found in compounds (XeF4, for example), they generally don’t like to associate themselves with other elements. Why is this? Well, let’s look at our electron configurations!
The general electron configuration for a noble gas is [core] ns2 np6, where “[core]” represents all electrons that aren’t in the valence shell and “n” represents an energy level (remember our quantum numbers?). For example, the electron configuration of neon (Ne) is 1s2 2s2 2p6, or, in the shorthand “[core]” notation, [core] 2s2 2p6.
(The reason we don’t use the standard [Noble Gas] notation is because, as our professor pointed out, this notation would have you think that, at higher energies, d electrons act as valence electrons, which isn’t true. In this notation, we include those in the core.)
As we can see from our electron configurations, noble gases have all of the eight valence electrons that they need to be happy: two in their outermost s orbital, and six in their outermost p orbital. This means that noble gases exist only as atomic allotropes, or allotropes that consist entirely of individual atoms interacting only through London forces.
Pretty simple, right? Personally, I find these dweebs pretty boring. If you move just one group to the left, however, you’ll run into the guys that happen to be my favorites…
Group 17/7A—The Halogens
A lot of times, when you ask someone to name a dangerous chemical, they’ll come up with some kind of halogen. This isn’t without foundation. Because halogens are situated at the near extreme of the periodic table (meaning they have a high effective nuclear charge) and lack but a single electron, they’re highly reactive when compared to other species. Among them is the most reactive nonmetal, my favorite (and environmentalists’ “favorite”), fluorine.
Halogens have a general electron configuration of [core] ns2 np5, which means that they only require one additional electron to be happy. This means that they very readily form diatomic molecules, combining into the very recognizable molecules F2, Cl2, Br2, and I2. The molecules in these diatoms are bonded to each other with single bonds, which is nice and simple. Unfortunately, it only gets more complicated from here.
Here resides every aerobic being’s favorite element, oxygen. (You’ll have to wait a minute to get to the one that all of us should share an affection for.) Being in Group 6A (or 16, if you have to be, you know, “right“ about it), it, along with all of its fellow group members, has a total of six valence electrons.
How does that translate into electron configurations? Well, the general configuration for members of this group is [core] ns2 np4. Practically, this means that an oxygen atom (or a sulfur atom, or a selenium atom) can’t single bond with another oxygen atom to get its full octet. Instead, complicated things occur.
Oxygen isn’t that hard to figure out, thankfully. You’re probably quite familiar with O2, a diatomic gas of oxygen that we usually just refer to as “oxygen.” The only difference between oxygen and chlorine here is that oxygen is bonded to itself with a double bond instead of a single bond.
(Another allotrope of oxygen, O3, called ozone, exists, but it’s not as stable, so I’m going to skim over it here. The Wiki, however, does a great job explaining what it is.)
However, when you go down to sulfur, things get a little trickier. As it turns out, sulfur isn’t a big fan of multiple bonds. The reason for this has to do with the jump in energy level from oxygen to sulfur: the orbitals participating in sigma (single) bonding in sulfur are larger, meaning that the orbitals that participate in pi (multiple) bonding don’t overlap as well. This ultimately results in S-S single bonds being more energetically favorable than S=S double bonds. Thus, while oxygen is over there making nice little diatomic gases, sulfur has to make big, bulky S8 crowns.
If you look at the molecule, you’ll notice that sulfur forms two single bonds with other sulfurs to get its full octet. It does this instead of forming the double bonds that oxygen does because, as we said before, it prefers single bonds to double bonds. This, turns out, is a rule for any element below the second row of the periodic table: stable pi bonding only occurs with 2p nonmetals.
And hey, speaking of multiple bonds…
This is where we find nitrogen, a fairly recognizable element that’s incredibly abundant in our atmosphere as its allotrope, N2. Since this group is the 5A group, we know that all of these elements have five valence electrons. Put in other terms, they’re all missing three valence electrons. Don’t believe me? Take a gander at the electron configurations, why don’t you?
(Ugh. I feel like I just won 500 Southern Grill points.)
The general electron configuration for the Group 5A elements is [core] ns2 np3. As you can see, we only have three electrons in our p orbital now, three less than our six electron maximum. While that does give us a nifty little half-full shell, our Group 5As are still trying their hardest to find three electrons.
Nitrogen does this by combining itself into its only allotrope, N2. This is a diatomic gas with, you guessed it, three bonds connecting the atoms together. This makes nitrogen gas freaking incredible in terms of stability, but it also brings a problem that we’d rather forget about to light: if nitrogen satisfies itself by triple bonding with itself, what does phosphorus, which hates multiple bonds, do?
Well, it turns out that there are a couple allotropes of phosphorus. The first, white (yellow) phosphorus, or P4, is a pyramid of four phosphorus atoms bonded to each other with three bonds each. In case you hadn’t guessed, none of the atoms are particularly happy in this configuration, since the configuration here is pretty strained. That means that P4 is reactive enough that it ignites instantly when exposed to air. Although this has given it numerous uses (it used to be used in matches, for example), it means that it is by no means the most stable form of phosphorus.
Another form is red phosphorus, which is an amorphous linear polymer with polyhedra throughout. There’s also black phosphorus, which can either be a covalent lattice or metallic material, depending on the handedness of the whole setup.
And speaking of black lattices that conduct electricity…
Here we have every living creature’s favorite element, the basis for all life, carbon. If you’ve taken organic chemistry at any level, you’ve probably had a professor who screamed at you, “Carbon is tetravalent carbon is tetravalent carbon is tetravalent you imbeciles.” If you haven’t, then, as you’ve deduced, carbon and its metalloid cousin, silicon, are tetravalent, meaning they have four valence electrons.
We can see this by looking at the electron configuration for Group 4A, which is [core] ns2 np2. Now we’re missing four electrons, which means that these guys have got some work to do if they’re going to be happy in life.
For carbon, this is relatively easy, at least in a limited sense. No longer able to bond with one other atom of itself to be happy, since quadruple bonds aren’t a thing (well, they are, but don’t tell organic chemists that), it has to settle for hooking up into large, interconnecting networks that we call covalent lattices. Both of the covalent lattices of carbon are common and extremely useful: they’re called diamond and graphite.
Diamond is a covalent lattice material wherein every carbon is bonded to four other carbons with single bonds. These are three-dimensional crystals, and constitute the hardest naturally occurring material (well, Theodore Gray might disagree with you). However, this isn’t the most stable allotrope of carbon. Graphite wins that award, with its two-dimensional alternating single and double bonds. (Yup, that’s right: your pencil lead is more stable than a diamond. This is the part where my professor suggests that we chemists start demanding graphite rings instead of diamond ones.) Graphite can conduct electricity in two dimensions, but since it consists of individual sheets that aren’t covalently bonded together, it doesn’t conduct electricity well in the third dimension.
There’s a third allotrope of carbon that was only discovered relatively recently, and that’s buckminsterfullerene. Also called “bucky ball,” it’s shaped like a soccer ball and has single-handedly ushered in a whole new branch of chemistry, but that’s a topic you just read a Wiki page about.
All right, so what about silicon? Well, it behaves essentially like diamond. That shouldn’t surprise you: 3rd row elements don’t like double bonds, remember?
Okay! You’re about ready to be done, right? I know I am!
This is where we reach the end of our allotropes. Boron, with three valence electrons ([core] ns2 np1) physically can’t make itself into compounds that satisfy its octet. Instead, it bunches up into little electron deficient clusters of ten atoms. If you’ve got a keen eye, you’ll notice that, on the periodic table, boron falls on that jagged line of metalloids, along with silicon (and arsenic and tellurium and…). This is the reason why.
Phew! We’re finally through! You never knew there was so much about allotropes that you didn’t want to know, right? Same. Same. Now that we’re through, though, we can explore some more interesting periodic trends, such as oxidation states in binary reactions.
Questions? Comments? Corrections? Throw them at me, please! Complaints? Flames? You can keep those.