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.

To everything there is a season

...and a time to every scientific topic under heaven. A time to blog, a time to refrain from blogging due to the constraints of a non-academic job, and a time to get one's schedule under control and resume blogging. Hello, and welcome back to Delusions of Grandeur! It's been a heck of a hiatus, but we're back to our regularly scheduled programming of a little of everything. There are many fine places I wish to take you, dear readers, from soil orders to paleoclimate, but let’s start with something that has been on my mind lately. Seasons.

For those of you reading from the tropics, I’d like to apologize in advance for the mid and polar latitude-centric nature of this post. Rest assured, I will devote a future post to an equatorial topic. For those of you reading from the mid-latitudes, in particular, have you ever noticed just how much of our culture(s) is informed by the reality of four distinct seasons? Think beyond the obvious. In nineteenth century Europe, the annual extravagance of the spring blossoms and autumn harvest had a clear impact on the Romantic movement, while the annual bloom of the cherry trees has long been important in Japanese culture. Moving away from flora, the sharp contrast between the summer and winter months in the mid-latitudes has inspired tales like how the hero Glooskap, in one of his many acts of magnanimity towards the Abenaki people, brought fixed seasons to the land. I could toss out some more examples, but the point is clear - away from the equator, seasonality shapes life. But what, exactly, shapes seasonality?

Don't worry if you missed it this year, you've got next season. I guarantee.

First, let’s define “season”. Technically, a season is any period of the year, as defined by weather or work. So, you can have a growing season, a dry season, a baseball season; all usages are technically correct. In fact, the word originally applied exclusively to the time when fields were sown. For the purposes of today’s investigation, we’ll be defining season as any one of the four approximately equal periods into which the solar year is divided, based on rotation around the Sun, and the tilt of Earth’s axis. At a future point, we’ll take a look at the phenomena of wet and dry seasons - those are governed by weather patterns, and as the names indicate, typically have to do with precipitation.

We’ll start with revolution around the Sun. 

 

As you all know, the Earth revolves around the Sun, with each revolution lasting 365 ¼ days. The orbit of the Earth is not quite even. Perihelion refers to the point in each orbital cycle when Earth is closest to the Sun, aphelion to the point when Earth is furthest. This plays a role in seasonality, but a very minor one. We need to look at axial tilt to get the full story.

So, axial tilt! What is it? In the words of many an English teacher, show, don’t tell.

The Earth spins as it rotates, with each spin lasting 24 hours. The Earth, however, is not quite spinning on a straight axis - the imaginary line one could trace from the North Pole through to the South Pole is not perfectly perpendicular to the ecliptic, the theoretical plane on which the Earth orbits the Sun; rather, the axis is tilted at 23.5 degrees relative to that the ecliptic.

Cool. What does this mean? Well, think about it. If the Earth’s axis was perpendicular to the ecliptic at all times, the Equator would always be closest to the Sun, the poles would always be furthest, and the mid-latitudes would always be somewhere in-between. Latitudes equally distant in opposite directions from the Equator (say, latitudes 45 N and 45 S, or the very approximate locations of Montreal, CA and Wellington, NZ) would consistently experience the same day length at the same time. As anyone from the Northern Hemisphere who has taken a winter vacation to the Southern Hemisphere (and vice versa) can tell you, this is manifestly not the case. 

Enter tilt. 23.5 degrees is just enough to ensure that, for portions of Earth’s revolution around the Sun, one hemisphere is closer to the Sun than the other. When one hemisphere is closer to the Sun, the days in that hemisphere are longer than the nights, and the amount of sunlight hitting the ground (and subsequently warming the air) is greater than at all other times. This is a consequence of Earth being round. Take a look at this image. 

The arrows represent incoming sunlight.

See how the sunlight hits Earth? Imagine a flashlight shining on a ball. At the “equator” of the ball, the light falls in one fairly tight beam. At the “poles” of the ball, the light is more diffuse (try it with some household objects, if you don’t believe me). The same factor is at play with the Earth. When a hemisphere is tilted towards the Sun, the incoming solar radiation in latitudes further from the Equator behaves a little more like solar radiation at the Equator. Got it?

Incoming radiation on the left is what you would get at the Equator, incoming radiation on the right is what you would get closer to the poles.

And there, in a nutshell, is the reason for the season. In the summer of either hemisphere, incoming solar radiation is greater, on account of tilt. For the same reason, winter is the result of tilt pointing the hemisphere away from the Sun. The effects are more pronounced at the higher latitudes, where the summers are characterized by constant sunlight, and the winters by a complete absence of sunlight (think about what parts of the globe are most tilted towards/away from the Sun, relatively). The orbit of the Earth around the Sun comes (slightly) into play here; perihelion occurs during the Southern Hemisphere summer, which means that this hemisphere receives more intense incoming solar radiation during its summer. As far as average temperatures go, however, there is very little difference between Southern and Northern summers. A greater proportion of the Southern Hemisphere is water, which is less sensitive to being warmed via solar radiation, and that evens out the variation.
 

Ok, we’ve got summer and winter explained; what about fall and spring? And what’s up with equinoxes and solstices?

We generally don’t celebrate solstices and equinoxes the way we once did, but there is a reason they are a feature of our calendars. The solstice, both summer and winter, refers to an extreme in the progression of Earth through the tilt cycle. Summer solstice, in a given hemisphere, is the point at which that hemisphere is most tilted towards the sun, winter solstice is the point at which that hemisphere is most tilted from the Sun. We’ve known about this since long before we had any ideas about the tilt cycle, because the solstices also correlated to day length. The summer solstice is the longest day of the year, the winter solstice is the shortest day of the year (note that “day” refers to length of the sun being above the horizon; the Earth always takes 24 hours to complete one revolution). 

Left-most sphere shows Earth at the Northern Hemisphere summer solstice, Jun 21 or 22

Hang on. Summer is marked as beginning at the summer solstice, winter as beginning at the winter solstice. The winter solstice is the shortest day of the year. Logically, that is the point at which the days start getting longer. But winter is cold. How can it be that the part of the year when day length is increasing is also the part of the year with, on average, the coldest temperatures?

Funny you should ask. One “Herb” asked the same question of a US Department of Energy “Ask a Scientist”, and received the following pithy reply from one Dr. Cook.

“It takes the Earth a long time to cool off. Heat has to be released from a significant storage of energy in the ground, trees, buildings, etc. While the cooling is taking place, a lot of energy is still being released into the atmosphere, slowing the cooling. So, the coldest temperatures come later, after this cooling has occurred (in January or early February).

The reverse situation occurs in the Summer. June 21 is not the hottest day of the year in the Northern Hemisphere, as it takes the a long time for the ground, trees, buildings, etc. to warm up; the warmer they become as the Summer progresses, the more heat is released into the atmosphere, resulting in the hottest days being later (July and August). “

Probably not taking place on June 21st.

Dr. Cook even answered the related question regarding the summer solstice. Thanks, Dr. Cook! Oh, and the technical term for this phenomenon is “thermal lag”. The period when there is less incoming sunlight, and the Earth is still cooling off covers the fall, and the period when there is more incoming sunlight, but the now-cooled Earth is still heating up covers the spring.

The equinoxes (vernal in the spring, autumnal in the fall) refer to the two points in the solar year when the Earth is actually parallel to the Sun. On these two days, day length is equal throughout the Earth. Fittingly, the equinoxes fall right between the solstices, as the solstices represent the two extremes in day length.

That’s seasonality, the short version. Now that we’ve covered this, we can devote later posts to interesting quirks of seasonality, like the important, but difficult to spell Milankovitch cycles. Or seasons on other planets; yes, other planets have seasons, and for more or less the same reasons. As for now, whatever hemisphere you’re in, and whatever seasonality you experience, get out there and enjoy it!

Coppery keys to chemistry

Here we are again, confronting another interesting chemistry topic. It’s back to the periodic table today, as we continue our march to the Nobel group. Before we go any further, I feel like I need to restate again what an incredible thing the periodic table is. As much as I joke about how dry the table can be, the fact that we have such a visual, user-friendly and logical method of ordering elements by physical and chemical characteristics might just be one of the most incredible accomplishments of chemistry. No matter how complex the chemistry, everything that happens is based on a relatively small number of laws, some general tendencies of matter, molecules and charges, and the characteristics of elements. The periodic table lays all of those characteristics out in a way that can be almost instantly grasped. No endless lines of text, no subdivided columns, cells and subscripts of subscripts. Can you imagine what a purely text periodic table would look like? Probably nothing good.

Tried to find an image, unable to find someone crazy enough to make an all-text table

 So, today, we’re going to be looking at that upper right section of the marvelous table. One of the fun things about moving towards the right side of the table is the increasingly cool names that the element groupings acquire - none of this “metalloid” business here (descriptive enough, but rather bland). We’re in to the vaguely alchemical-sounding Greek-derived roots here, and today’s sterling example of nifty names in science is the chalcogens.

 

Chalcogen, for those of you not as up on your ancient Greek, translates (roughly) to “copper former”. In addition to the chemical characteristics we’re about to explore, the chalcogens all tend to be found in copper ores. Keep the “-gens” ending in mind, as we’ll be encountering it again next week.

 

They're in here somewhere...

So, besides hanging out with copper, what do the chalcogens have in common? First of all, they are not metals. Additionally, oxygen, sulfur, and selenium all have 4 electrons (out of a maximum of 6) in the outermost valence shell (as does tellurium, but it is counted as a metalloid). Remembering back to what we’ve seen of how atoms achieve that coveted full outermost valence shell, a mostly-full valence shell is much more likely to gain additional electrons than it is to shed electrons.

Chalcogens can either gain two electrons and become an anion with a charge of -2, or else form covalent bonds (typically with other non-metals). In fact, covalent bonds involving chalcogens are probably the single most important bit of chemistry on the face of the Earth. Remember back when we took a look at hydrogen bonding, that weird form of bonding that was only possible because the covalent bonds between oxygen and hydrogen atoms in water molecules contain some charge? Those semi-charged bonds are a direct consequence of the fact that water, as a chalcogen, is more likely to gain an electron than to lose it. Even when in a covalent bond, electrons are going to be drawn to the chalcogen. And that is how we wind up with the partially charged covalent bonds that make hydrogen bonds possible. 

 

Thank you, electron-gaining tendency of chalcogens.

Oxygen is easily the most recognizable element within the chalcogens - the name recognition may have something to do with the fact that we breathe oxygen. Without turning this into a post about human physiology, I’ll say that we, along with quite a few life forms on Earth, use oxygen at a cellular level in order to convert sugar molecules into energy through aerobic respiration. Oxygen is quite reactive, and capable of combining with just about anything. Oxygen is also abundant - it makes up 21% of the atmosphere. This was not always the case. Way back in the Archaean, billions of years ago, there was relatively little oxygen present in the atmosphere, and those organisms in existance practiced anaerobic (no-oxygen) respiration. However, oxygen is by-product of photosynthesis. The rise of photosynthetic organisms raised the oxygen content of the atmosphere, which, in turn, made it possible for more organisms to use the more efficient aerobic respiration, leading to life as we know it.

Not pictured: life as we know it.

Enough about oxygen for now. How about sulfur? Well, it’s also fairly common. Sulfur is the tenth-most abundant element in the universe, and crops up with some regularity in both mineral compounds and fossil fuels. The later can be an issue, as burning fossil fuels releases to the atmosphere, among other things, sulfur. Once there, the sulfur forms acid rain, with corresponding consequences for any ecosystems receiving that low pH precipitation.

How did sulfur wind up in fossil fuels in the first place? Aren’t those largely hydrocarbons? Well, yes, mostly. Fossil fuels are the fossilized remains of plants and animals, which, like most life on Earth, are mostly carbon, hydrogen, oxygen and nitrogen. Those four elements can be expanded out by the addition of sulfur and phosphorus to a group of six elements, which make up the vast bulk of all organic molecules. So, those fossils in the fuel may well have contained sulfur prior to being fossilized.

Alternately, sulfur came into the picture a little bit later. Sulfur belongs to the same group as oxygen, which would imply that sulfur behaves at least somewhat similarly to oxygen. And, in fact, there is an entire class of bacteria that live in non-oxygenated environments, and use sulfur to fulfill much the same role as oxygen in respiration. The process of converting organic matter into fossil fuel commonly involves these sulfur-oxidizing bacteria, leading to sulfur being incorporated in fossil fuels, and released upon burning.

It's round, it's pink, it breathes sulfur. Cool!

How about selenium? Well, selenium is the most metallic of the chalcogens, to the extent that it is occasionally included in the metalloids, and can be used as a semi-conductor. Selenium occurs in minerals (including that aforementioned association with copper ores), and is a nutrient to plants and animals, in small doses. Selenium is toxic at high doses - in certain soils, selenium is highly available, and accumulates in plants. When cattle eat those plants, they develop selenium poisoning - a fun fact admittedly not overly connected to the chalcogenic properties of selenium, but of note to any readers thinking of starting a cattle farm. Test your soils for selenium!

The chalcogens are a small group, but what they lack in numbers, they make up for in how many roles they play in chemistry. It figures that, with elements going up to number 118 (and rising), you're going to get a certain number of overachievers like oxygen. The nonmetals are a real cluster of such overachievers, as you might guess by the fact that we had to separate seven elements over two weeks, when the 38 transition metals all fit neatly into one post. It may be the natural bias towards organic chemistry of an organic life form, but I would argue that much of chemistry is contained within this small group of elements, together with next week's subject, the halogens. Stay tuned!

Bomb(ardier)s away!

The astute of you may have noticed that we’re running a little bit behind here, post-wise. So, I’m going to switch things around a little bit, and give us another break from the periodic table. It’s still Chemistry Wednesday, but we’re going to take a look at another piece of applied chemistry, this time, chemical ecology. Or chemical entomology, if you want to call it that. Or exploding beetles of doom, if you want to be awesome and call it that.

Definitely exploding beetles of doom.

Yes, the bombardier beetle, an organism so bizarre that its very existence is a fundamental part of the intelligent design argument - the beetle displays such an off-the-wall defense mechanism that it couldn’t possibly have evolved. Well, actually, it could have, and it did, and we’ll see why and how in a moment. But first, the chemistry part. And, in the interests of full disclosure, I must say that the guy who discovered all of this was one of my favorite professors in college. Hats off, Dr. Eisner.

Well, actually, first a crash course in the beetles. Bombardier beetles belong to the Carabidae family. These are ground beetles - they’re quick and nimble scuttlers, but not particularly strong or fast flyers. When Carabidae beetles get in trouble, they need a response other than taking wing, and boy, do bombardiers ever have a response.

Imagine that you are an ant. There, on the ground ahead of you, is a large, potentially succulent bombardier beetle. It’s got small, tightly furled wings - it’s going to be a ponderous process of getting airborne, and you’re a quick little thing. So, you run at the beetle. Bang! You’ve suddenly been hit with an extremely well-targeted, 100 degree Centigrade, chemically irritating fluid. What just happened?

The ant's eye view of what just happened.

Say hello to benzoquinones! And some other things. Benzoquinone, specifically 1,4 benzoquinone, is an extremely strong irritant, one that can cause blisters, blindness and respiratory damage in humans (and most other life forms). Clearly, being able to squirt this stuff at an attacker is a good survival strategy. It’s like taking the painful, itchy experience of encountering a poison sumac, cubing the pain, and subjecting anything that gets near the tree, not just touches it. Bombardier beetles aren’t overly fond of benzoquinone, mind. Part of the precision inherent in spraying is to allow the beetles to avoid accidentally spraying themselves.

Bombardiers avoid any negative effects of carrying this stuff around inside of them by, well, not carrying it around inside them. Instead, bombardiers carry precursors to benzoquinones, and combine these precursors when threatened, in order to synthesize the irritant spray. Bombardiers have a specialized arrangement of glands at the rear of the body - two glands empty into a reaction chamber, which then empties into a firing mechanism. The reaction chamber, as you might expect, is made of rather resistant tissue.

With a name like "explosion chamber", it would almost have to be.

One of the two glands contains hydroquinones (a chemical that can be an irritant in sufficiently high doses, and is involved in regulating melanin) and hydrogen peroxide, which appears in everything from topical antiseptics to rocket fuel. The other chamber contains a mixture of enzymes. When the two glands are emptied into the reaction chamber, enzymes break hydrogen peroxide down into hydrogen and oxygen. The oxygen then converts the hydroquinones to quinones by removing two hydrogen atoms from the hydroquinone molecule.

Not pictured: heat

Ok, that explains the synthesis of a toxic spray, but not the explosive burning part. Ah, but it does. Remember that bit where hydrogen peroxide was split in half, generating oxygen? As we’ll see next week, oxygen is a very reactive element. That’s where the explosion comes from. The heat is the result of the reaction being exothermic. Exothermic reactions release heat. Sometimes, the heat is barely perceptible, and sometimes it heats the resultant solution to the boiling point. Really, it’s just a matter of magnitude. And for the record, the reaction is controlled, so there’s no risk of the beetle exploding - the reaction chamber can withstand the force of the explosion, particularly since said explosion is directed outward.

Now, we’ve taken a look at how the chemistry works. How on Earth did all of this evolve? Well, as tends to be the case in evolution, bombardier beetle spray arose through a sequence of relatively minor tweaks to existing mechanisms. Let’s run this backwards.

There is a beetle in the Carabidae family that synthesizes the same irritating mix of quinones, but rather than shooting them out at high temperature, exudes them as a high-temperature foam which bubbles off the resistant carapace of the beetle (the foaming action should be familiar to anyone who has ever used hydrogen peroxide to clean out a cut). The reaction is still exothermic and explosive, but it does not take place within a reaction chamber that empties into a firing mechanism. You might see a flaw with this method, however. By exuding the irritant as a foam, the beetle can only deter predators that have made it onto the body - a really determined predatory might be able to hang on. Beetles which could spray the exudation would have a higher chance of surviving encounters with a predator. It’s a relatively minor evolutionary matter for the reaction chamber to become more rigid, and for an aiming mechanism to develop.

 

It's the Windows 98 of bombardiers. Almost there.

All right, but how do we get to the hot, toxic exudation in the first place? Well, any number of insects secrete hydroquinones as defensive compounds. Most insects also produce small amounts of hydrogen peroxide as a byproduct of metabolic reactions. Over time, a beetle which initially secreted hydroquinone via a gland and duct arrangement may have begun to combine hydrogen peroxide with hydroquinone, generating both a more toxic secretion, and a secretion with more temperature oomph. It’s easy to see the evolutionary progression from a beetle which exudes a hydroquinone solution to a beetle which combines hydroquinone with hydrogen peroxide before exuding it.

Hydroquinone, in turn, is a tweak on a more primitive defensive compound, quinone. Quinone was not originally a defensive compound - it had other, structural roles to play in insects, but eventually began to be used as defensive secretions, as they have an unpleasant taste. What was that original structural role? That would be protective pigmentation - various quinone derivatives are widely used in order to provide some protection from solar radiation. You might be familiar with one of those derivatives - melanin.

So, however exotic the chemistry of the bombardier beetle may be, just remember, it’s just an evolutionary hop, skip and jump away from human skin tone. That’s wicked cool.