The fantastic four

Wow, can you believe that we’re nearly through the periodic table? I believe it was way back in March when I got the crazy idea to work through the entire thing. Fear not, Chemistry Wednesdays will continue, but the topics will get just a little bit more unpredictable. But, before we reach that point, we’ve got three more groups of elements to go through. We’ve looked at the metals, and the semi-metals. This can only mean one thing is coming up - the non-metals!

 

Helpfully marked in orange

Yes, the non-metals, a group of elements largely defined by what they are not. Also, a somewhat misleading name, as all non-metal elements are not metallic,but not all non-metallic elements are referred to as non-metals. Confused yet?

All right, remember in the good old days of the alkali earth metals, when everything was simple, and each group of elements simply consisted of one of those straight-down-the-table groups like Group I and Group II? Then along came the transition metals and screwed all of that up, by being very similar to one another across groups. As we move towards the right side of the table, groups become more distinctive relative to one another once again. The last two groups on the table are usually examined as separate categories, as is (sometimes) the third-to-last. We’ll look at those next week, as I don’t fancy covering carbon chemistry and oxygen chemistry in one post. Today, we’re going to look at a kind of catch-all group - those non-metallic elements that do not belong to one of the three distinctive final groups on the table. Don’t worry! There are only four of them!

Most of the elements in this group are clustered together. Most, but not all. Yes, today is the today when we finally look at hydrogen which, in spite of hanging out on top of Group I winds up being discussed along with carbon, oxygen (sometimes) and nitrogen (that little tetralogy of elements is the basis for an entire branch of chemistry, which we will get to in a moment).

Ok, so besides being excluded from the metal club, what exactly distinguishes non-metals from the pack? Well, they lack the properties of metals - ductility, ability to form alloys, tendency to efficiently conduct heat and electricity, etc. Non-metals also have more electrons in the valence shell, usually from four to eight electrons (remember the oddities of the d and f orbitals, which keep them from ever being proper valence shells). The non-metals we’re looking at today usually have four or five valence electrons (hydrogen, with only one electron, being the exception). With half-filled (or more) valence shells, the non-metals are more likely to bond in such a way as to fill the valence shell. Contrast this with the tendency of metals to shed valence electrons.

Non-metals, then, tend to either be the negative side of an ionic bond, or bond covalently. The elements we’re looking at today are all more likely to bond covalently - we’ll get to the ionic bonds next week. This tendency to form covalent bonds gives rise to a rather important bit of chemistry - organic chemistry.

Organic chemistry is defined as the branch of chemistry concerning carbon compounds. After carbon, the main elements involved in organic chemistry are hydrogen, oxygen and nitrogen, following the easy-to-remember acronym “CHON”. Phosphorus and sulfur are two common additional guests, and, not coincidentally, also in the non-metals group (although, again, oxygen and sulfur are sometimes separated out).

Organic chemistry is sometimes called the chemistry of life - it deals with the compounds that form all living tissue, from DNA bases to skin and sinew. Certain characteristics of the nonmetals, particularly carbon, are responsible for the primacy of organic chemistry in biology, so let’s focus on carbon for a little bit.

 

Thanks, organic chemistry!

Carbon (abbreviated C) is interesting, in that it is capable of covalently bonding with four other atoms at once, due to its 4 valence electrons (in an 8-max shell). Silicon, germanium and tin are also (theoretically) able to do this, but, to varying extents, all three of those elements are somewhat likely to give up some or all of their valence electrons, due in part to the presence of a d-orbital. Carbon is the only 4-valence electron element with no tendency to give up electrons, and thus almost always forms covalent bonds. Carbon is also the sixth-most abundant element in the universe. Combine these two things, and the fact that life on Earth tends to be carbon-based starts to make sense. Covalent bonds, after all, are possible between elements of all electron configurations (in contrast to metallic bonding), form easily at the standard range of temperatures on Earth (in contrast to a form of bonding we’ll look at in a few weeks), and don’t dissolve in water (in contrast to ionic bonding). That’s reassuring, isn’t it?

Nitrogen (N) has interesting bonding abilities, as well. Nitrogen has one more valence electron than carbon, and needs three additional electrons to form a full shell. Now, plenty of N atoms achieve this by bonding to three additional electrons, but some N atoms achieve this by forming a triple covalent bond to another N atom. In two triply bonded atoms, three pairs of electrons are shared, forming a very strong bond.Theoretically, any atom with the right range of electrons in the valence shell can form double or triple bonds (C, ever the overachiever, is capable of forming a triple bond with another C or N atom, and a single covalent bond with another atom, putting a certain amount of pizzazz into organic chemistry), but N is especially likely to form triple bonds. Two triply bonded N atoms form a stable molecule, as both have full valence shells. In its stable, 2-atom molecule, N exists as a gas. And not just any gas. Gaseous N makes up 78% of Earth’s atmosphere. Nitrogen is the fifth most abundant element in the universe, and plays all kinds of interesting roles in ecology. Which, one of these days, I will get into.

Phosphorus (P) behaves somewhat similarly to N, but a) the bonds aren’t as strong, c) there is a d orbital involved, and c) P has a tendency to spontaneously combust in air. Let’s make a very long story short and simply say that P has a very strong affinity for oxygen, and rapidly reacts with any available oxygen to form certain compounds. As you may remember from Group I chemistry, rapid reactions sometimes generate, well, fire. Who says organic chemistry is boring? Phosphorus is an important nutrient, and an ingredient in explosives. How many elements can you say that about?

That leaves hydrogen, our weird little outlier. Now, the reason we didn’t stick hydrogen in with Group 1 is that, in spite of having only one electron, hydrogen still has a half-full valence shell. So while hydrogen can lose that one electron, it’s just as likely to acquire another electron via covalent bonding. Hydrogen may exist as a gas (two hydrogen atoms can bond covalently forming a stable molecule), or in combination with another element - in fact, hydrogen naturally combines with every element in the periodic table, except for the ones at the bottom (they don’t exist in nature) and the Nobel gases (wait a few weeks for the answer to that one). Hydrogen is the most abundant element in the universe, comprising something like 90% of all matter. Hydrogen pops up everywhere from organic chemistry to fertilizers, to the hearts of stars, where it burns in a process that, if only we could understand exactly how it works, could fuel the future of humanity. 

Yeah, I'm a playa

Wow. That was a lot of information for four elements. Non-metals are quite the grab bag, eh?