Noblesse oblige

It's Wednesday, and you all know what that means. Chemistry! We've still got a corner or two of the periodic table left to visit, and today, we'll be taking a trip to the right-most column of the table, where the noble gases hang out.

Hi

So! What is special about this group of elements? Are they explosive? Poisonous? Metalloid? No, no, and no. The distinguishing characteristic of the noble group is more what they do not do. As we've worked our way across the periodic table, I'm sure you've noticed that elements are almost never found as a group of unconnected atoms. They may bond in pairs of atoms to form a gas, like nitrogen, or perhaps engage in the complex bonding we see in the transition metals. Maybe they form covalent bonds, maybe they form ionic bonds. Perhaps they bond to other atoms of the same element, perhaps they join with atoms of other elements to form compounds, but just about everything out there tends to form bonds. Except for the noble gases.

Weirdos

This is a concept that has come up repeatedly over the course of Chemistry Wednesdays, but here it is again; valence shells. Remember, all of the varieties of chemical bonds we've seen have involved that outermost theoretical shell of electrons. Atoms combine with other atoms in order to achieve a full outermost shell, or orbital. You may have noticed a trend as we have moved across the table. Way back on the left side, in the alkali metals, that outermost shell had only one electron. By the time we got to the halogens, that outermost shell was only one electron short of a full set. That trend is real, and here, at the right side of the table, we reach atoms that naturally have a full valence shell.

Atoms don't look like this in real life, but let's pretend.

If you're thinking that's a little odd, you are not alone. In fact, the periodic table, as originally conceived by Mendeleev, did not allow for such a thing. However, in 1895, two British chemists discovered argon; while they initially wondered if it wasn't simply a variation of nitrogen, it quickly became clear that there was a whole new family of elements out there to be discovered. The ability of the noble gases to exist in an unbonded (or monatomic) state was one of the things that pointed early 20th century chemists towards a more accurate understanding of electron bonding. 

As a group, then, the noble gases are incredibly unreactive. This was a contributing factor in why it took us so long to learn of their existence. Now, just because the group is unreactive does not mean it never reacts. No compounds have been made containing helium or neon, but argon, krypton, and xenon compounds have been created, and have a few chemical and industrial uses. For the most part, though, we're most likely to encounter them in their monatomic, gaseous forms. 

And encounter them we do! Let's start at the top, with helium. Helium is the second-most abundant element in the universe (after hydrogen); it is a major component of the Sun, and is present in trace quantities in Earth's atmosphere. Pure helium gas has a density about one-seventh that of air on Earth, and is used in lighter-than-air craft, from hot air balloons to party balloons. Helium has so many applications (everything from MRIs to silicon computer chip manufacture) that we're actually facing a helium shortage. Well, high demand, limited availability on Earth, and some odd quirks of government policy.

Brought to you by helium. And a tremendous yardage of nylon.

Neon might be most familiar from the eponymous lights. Actually, all the noble gases can be used to create the so-called gas discharge decorative lighting, which is commonly called neon lighting. In this form of lighting, a noble gas is exposed to electrical energy. The energy "excites" the electrons; they briefly jump to a theoretical higher energy level, then return to their original energy level, and emit a colorful glow in the process. Neon is also used as a coolant, and in the manufacture of lasers.

And yes, these are the appropriate colors for each element.

Argon is the third most abundant gas in Earth's atmosphere (after nitrogen and oxygen). There are a number of industrial uses for argon, mostly stemming from its lack of reactivity; argon gas is used to buffer reactions and equipment that might be damaged by interaction with other, more reactive gases (I'm looking at you, oxygen). 

Other uses include lovely blue and violet lights.

Krypton, in addition to being Superman's home planet, has a few uses in flash photography, and is used in some fluorescent light bulbs. Krypton is difficult to obtain, which has limited its industrial uses.

Interestingly enough, krypton, unlike Kryptonite, is not a sickly greenish color.

Xenon turns up in strobe lighting, and car headlights. Certain kinds of xenon lamps emit a light strong enough to kill bacteria, and xenon has been used as a component of experimental rocket fuel.

Xenon in action.

Radon is the only radioactive noble gas. It forms as a component of uranium decay; since uranium is present throughout the Earth's crust, radon can and does form anywhere. Depending on the underlying soil and rocks of a given area, radon can accumulate in homes, where the radioactive gas is inhaled, to the detriment of the one inhaling. Radon is the leading cause of lung cancer amongst non-smokers. It is recommended that people living in areas prone to radon accumulation regularly test their homes and workplaces for the gas. The month of January is National Radon Action Month in the United States, and even if you don't live in the US, it's good to make sure that radon is not an issue in any buildings where you might spend a lot of time!

It's serious stuff

There's the noble gases for you. From hot air balloons to radioactive gas, they get a lot done for a group that was once called the "inert gases". Next week, we'll finish out the table with a look at those weird "uup" and "uus" elements you see hanging out in the bottom right row. It's been a long ride; thanks for sticking around!

Let's start off 2014 with some chemistry

Happy New Year! In addition to being the first day of solar year, as marked on the Gregorian calendar, today has the distinction of being a Wednesday. You may all remember what that means. It’s time for another Chemistry Wednesday!

Our march across the periodic table came to an ignominious pause a year and a half ago with the chalcogens, but we’re going to dust ourselves off and proceed to the halogens.

As is our wont here at Delusions, we’ll start this sojourn into chemistry with a detour into etymology. Halogen is a relatively new word; it was coined in 1842 by a Swedish chemist from the Greek words for “salt” and “producing”. Right off the bat, this clues us in to the halogen group being somehow involved in the chemistry of salts.

Hang on, perhaps we should have started with a working definition of salt.

It's a start...

In chemistry “salt” simply refers to a class of compounds that are formed when an acid and base react and contribute ions (charged particles) to a reaction. The acid contributes a cation (positive charge) and the base contributes an anion (negative charge); the anion and the cation then bond ionically (negative-positive). The resulting molecule looks a bit like this.

The anions are the blue circles, cations are the red circles. Note the regular form of the compounds; salts form crystals.

A common salt is NaCl, better known as table salt. Sodium (Na) is the cation in the mix. We’ve met sodium before on this blog, and seen how it can bond with anions, like chlorine. Hey, guess what group chlorine (Cl) belongs to!

The halogen group, as far as chemical behavior goes, can be thought of as the opposite of the alkali metals. While alkali metals have only one electron in their outermost valence shell, and best reach a stable chemical state by shedding that electron (and becoming a cation in the process), halogens are one electron short of a full outermost valence shell. The best way for a halogen atom to reach a stable chemical state is to gain that additional “missing” electron. I think you can see where this is going.

When not combining into salts with alkali metals (or other metals), halogens in nature are found combining with other elements. Halogens are some of the most reactive elements in the periodic table, with fluorine being the most reactive of all the elements. Halogens are never found in an uncombined state in nature; that almost-full outermost valence guarantees some sort of reaction for any individual halogen atom.

Let’s meet the elements involved on a more detailed basis. In addition to being the most reactive of the elements, fluorine, like all the halogens, is poisonous. Fluorine exists in an array of compounds in nature, and is, in fact, the most abundant of the halogens. Fluorine that is bonded to another fluorine exists as a pale yellow (and extremely toxic) gas at room temperature, but fluorine compounds, such as NaF, are useful in the prevention of tooth decay, hence the addition of ionic fluorine, or fluoride, to many municipal water supplies. This is not the last time we will encounter a toxic halogen bonded to an explosive alkali metal, yielding a harmless compound. NaCl follows the same pattern. Other fluorine compounds turn up in everything from glass etching to refrigerator coolant.

Fluorine, in its gaseous state

Chlorine is also gaseous at room temperature, although it is more green than yellow; in fact, the name comes from the Greek word for light green. While slightly less reactive than fluorine, chlorine is still a very reactive element, and has some decidedly unpleasant applications; the gas was first weaponized in World War I, and has since been used to deadly effect in conflicts as recent as the Iraq War. More benign uses of chlorine include bleach, treating drinking water, swimming pool disinfectant, and industrial processes ranging from fabric dyes to plastics manufacture. Chlorine is a component of table salt, and of several other salts found in ocean water, to the extent that chlorine ions (chloride) are one of the most abundant present in ocean water.

This might be called too much of a good thing.

Moving down the group, we encounter bromine, which is a liquid at room temperature. In addition to being extremely poisonous, like the rest of the group, bromine is so smelly that its name comes from the Greek word for “stench”. Bromine was once used in the production of leaded gasoline; with that no longer being produced, most bromine is now used in photography. Bromine, as part of a compound secreted by an ocean mollusc in the Murex genus, was an integral part of the production of Tyrian Purple, one of the most renowned dyes of the ancient world.

Meet the Murex. Interestingly, the process of making Tyrian Purple cloth was notably smelly, likely due in part to the bromine component.

Iodine, contrary to how many of us may have encountered it at the drugstore, is a solid at room temperature. It is also naturally violet. Iodine occurs in ocean water and seaweed, and while it is toxic in large doses, at trace quantities, it is an essential nutrient for humans. Iodine deficiency can cause a swelling of the thyroid gland, also called goiter. The condition is now rare in the developed world, due the the widespread practice of adding small quantities of iodine to table salt (this is what “iodized salt” means), but goiter remains widespread in other parts of the world, mostly as a result of dietary limitations. Other health applications of iodine include tinctured iodine, or the brown bottled topical disinfectant you may have lurking in your medicine cabinet. This is made by suspending iodine in alcohol. Iodine (along with bromine) turns up in halogen lamps, and has a few applications in film processing.

Last, heaviest, and least reactive of the halogens is astatine. Astatine is extremely rare, and was only discovered in 1940. It is also a solid at room temperature, but is difficult to study; astatine is highly radioactive, and turns into the elements bismuth or polonium in a matter of hours. In fact, the name for astatine derives from the Greek word for unstable. The only real applications for this element are laboratory studies of radioactivity.

So there we have the halogens. Many of them are quite dangerous, but used properly, can be tremendously beneficial (and in some cases essential) to life. Chlorine can be a killer, but as component of drinking water treatment, has saved many more lives than the weaponized form has ended. Fluorine and iodine can be nasty in large doses, but trace doses keep our teeth intact and our thyroids functioning normally. The halogen group certainly illustrates how a little bit of information can be a dangerous thing, but at the same time, reminds us that science is neither inherently good nor evil. Research shows us how things work, it’s up to us to determine how we want to use that knowledge.