Using techniques pioneered by IBM Research Zurich, this is the first ever AFM image of a single molecule. The most detailed single image of a pentacene molecule you'll ever see, in fact.

Seriously, I'm amazed. I had no idea atomic force microscopy could be this powerful. The researchers have apparently achieved this stunning resolution by coating the tip of their microscope probe with carbon monoxide. I'm a little hazy on the details, to be honest, as I haven't had a chance to read much about it just yet.

The most amazing part if that you can actually see the chemical bonds in the molecule, as regions of electron density. You can even see the bonds to the peripheral hydrogen atoms. Interestingly, there seems to be a lot more electron density on the two end rings of the molecule. This seems to fit quite well with what we know about pentacene. As aromatic molecules go, it isn't really very aromatic at all. Its electrons aren't as delocalised as many others. As a result, chemical reactants tend to attack the centre of the molecule, where the bonds are weakest.


Source: BBC Science and Environment

Well that's interesting... A while back, I wrote a post, How to quench your thirst on the Moon, about how to find water for a lunar base, in which I mentioned that no one had really considered obtaining the water chemically from Moon rock. As it happens, I wasn't entirely correct.

About a month later, Nature News published an article entitled How to breathe on the Moon -- about chemically reacting oxygen out of Moon rock. So someone out there clearly is thinking along the same lines. Yay for chemistry!

This looks rather promising. I might have to try and get a copy of the paper and have a closer look!

Every now and again when I'm studying I discover something which I think is just cool. And relativistic quantum chemistry -- is just cool! Quantum mechanics and general relativity aren't the best of friends. This much is pretty well known, and a great many people (physicists and otherwise) have waxed lyrical about it. Finding a way of combining the effects of the two into a single theory is the goal of any theorist working towards the so-called "theory of everything." General relativity has difficulty fitting into the world of particles. Special relativity, on the other hand, bears no such hindrance.

In fact, special relativity underpins a huge amount of modern physics. It shouldn't come as any surprise, then, that relativity rears its head in physical chemistry too. A lot of chemists may not need to care about relativistic effects, particularly organic chemists. It only really becomes an issue with the heavyweights near the bottom of the periodic table.

Dating back to Bertha Swirles' 1935 paper, "The relativistic self-consistent field," the idea is quite intuitive. Special relativity asserts that the faster an object travels, the greater its effective mass. The key to all of this is in the ever-curious fine structure constant, α;



α is a ubiquitous little thing. It seems to crop up all over physics, but in this case it denotes something quite simple -- the speed of an electron orbiting a hydrogen nucleus. Hydrogen is the simplest atom. One proton and one electron. In a hydrogen nucleus, an electron will travel at roughly 137th the speed of light.

Increase the size of the atomic nucleus though, and you increase its charge. The increase in charge causes the electrons to accelerate to increasingly higher speeds. Nuclear charge is usually denoted by a Z, and it's blissfully easy to plug this into the equation. Just multiply by Z to find the speed of the innermost electron in the shell. Like so;



For instance, that atomic nucleus up at the top of this post is uranium. Everyone's favourite actinide. Uranium has 92 protons and thus, a charge of 92 on its nucleus. As a result, its innermost electrons will be travelling at around 67% of the speed of light. This increases their mass by around 34%, causing those electrons to orbit sligtly closer to the atomic nucleus. In turn, this has a knock-on effect, causing the entire atom to shrink slightly!

Interestingly enough, a whole range of effects are opened up by this relativistic malarkey. It contributes to many of the heavier elements having smaller atomic nuclei than they should. It explains why lead doesn't have the same crystal structure as diamond. It explains why gold is actually gold! The electronic energy levels in gold are shifted by the relativistic effect, causing it to absorb the right optical frequencies to make it appear golden.

Most dramatically, it explains why mercury is liquid at room temperature. Relativistic effects cause the contraction of mercury's 6s² orbital, severely weakening its ability to form bonds. Subsequently, mercury remains liquid down to nearly -39°C. When it boils at around 357°C, it's atoms drift away independently as a monatomic gas. Gaseous mercury is thus sometimes referred to as a "pseudo-noble gas."

Ahhh. I love finding things out.




...wait, what do you mean quantum chemistry isn't cool?

As any schoolkid who's into pyrotechnics will tell you, mixing an oxidising agent with any form of carbohydrate will usually get you a dramatic reaction. Trust me. I used to be one of those schoolkids. Several of the so-called "explosives" in the infamous Jolly Roger's Cookbook, like glycerine with potassium permanganate, were based on this same idea. Reactions of this type are surprisingly potent. So potent, you can use them to power model rockets!*



This classic YouTube video shows much the same kind of reaction. A gummi bear, made primarily of sucrose sugar is dropped into molten potassium chlorate. Fwoosh! Violent oxidation reaction! Sucrose, in this case, is a partially oxidised hydrocarbon. Stable under normal conditions, but an oxidiser can easily oxidise that sugar fully into water and carbon dioxide, which is what's happening in this video.** The byproduct from the oxidiser itself is potassium chloride, commonly found in low-sodium table salt.

This same trick can be used to make the deliciously named rocket candy!

Rocket candy is a really simple type of solid rocket propellant. The oxidiser used is potassium nitrate (KNO₃), which is a rather milder oxidiser than potassium chlorate! It's also often found in fertillisers, making it another old Jolly Roger favourite. To make it into rocket candy, you need to first melt a quantity of sugar. This needs to be done using an oil bath or similar, to prevent any hotspots which could cause the stuff to prematurely ignite. Once the sugar's all melted, you can simply mix in the KNO₃, pour it into the rocket casing and leave to solidify. Voila! A bit like Brighton rock, only a lot more flammable!

I'd imagine professionally made ones are probably made with powders which are sintered together at low temperatures, though trying to do that at home is a lot more hazardous. It's perhaps not the best of ideas to go putting rocket fuel in an oven...

Rocket candy works so well, it's the propellant of choice for most amateur rocketeers. As such, a recently announced project called Sugar Shot to Space aims to get the first amateur rocket into space using only an amateur propellant -- as opposed to the more high end propellants such as ammonium perchlorate used professionally. I certainly plan to keep an eye on their developments. It'll be intriguing to see how well they do!



*Yes, yes. I used to play with model rockets too. What can I say? Rockets are fun!

**Incidentally, don't try this with any old hydrocarbon. With things like fuels which aren't partially oxidised, the results genuinely can be explosive -- which is why MSDS sheets always say to keep fuels away from oxidisers!

I was flicking through the rather marvellous webcomic, Abstruse Goose, recently, when I happened upon this particular comic (click the panel for the full version). Wait, what? Chlorinated sugar? This piqued my curiosity somewhat, so cat-like as ever*, I decided to investigate further. It turns out that, as they say, many a true word is spoken in jest.

Make no mistake. The artificial sweetener sucralose, more popularly known as Splenda®, is precisely that. Chlorinated sugar. It's also known as 1',4,6'-Trichlorosucrose, E955 or (deep breath!) 1,6-dichloro-1, 6-dideoxy-β-D-fructofuranosyl-4-chloro-4-deoxy-α-D-galactopyranoside. Containing chlorine atoms covalently bonded to carbon atoms, it's an organochloride compound (also known as a chlorocarbon). This makes me raise a slightly skeptical eyebrow. Now, I have nothing against chlorine, but the releasing of chlorocarbons into the environment on a large scale? That makes me a bit uneasy.

I'll make no secret of the fact that I detest artificial sweeteners. As opposed to the good old fashioned sugar that our digestive systems have spent hundreds of millions of years evolving to digest efficiently, many people seem to favour fabricated chemicals which apparently taste "sweeter". As far as I'm concerned, they spoil the taste of things. Since when did "no added sugar" come to mean "icky artificial sweeteners added instead of sugar"? Couldn't they just - you know - not add any sugar? But enough of my ranting. I have a little more than just personal opinion to go on here.

The two molecules to the left are sucrose and sucralose. A disaccharide, sucrose is made from one molecule of fructose and one of glucose. It's the dominant type of sugar in sugar cane and sugar beet. Just plain ordinary table sugar. The same kind found across the world.

Sucralose, as the Splenda® people rightly claim, is "made from sugar". Just chop off a couple of hydroxy groups and swap them with chlorine (which is, admittedly, easier said than chemically reacted). In fairness, it is perhaps the least awful tasting of the artificial sweeteners still available**. In fact it seems to be steadily replacing aspartame, which is arguably a good thing.

So what's the big deal? What's my problem with chlorinated sugar? The trouble with organochlorides is that, while not necessarily toxic, they're notoriously stable. A good example would be a certain famous chlorocarbon.

Now let me make it quite clear -- I'm not comparing sucralose to DDT. They're very different chemicals, with different uses, reactivities, stabilities and so forth. The point to this paragraph is more to highlight precisely how stable large chlorocarbons can be, with DDT being a prominent example. DDT was used widely as a pesticide during 1940s and 1950s, due to it being extremely toxic to invertebrates, but "safe" to vertebrates like people. Except that it isn't as safe as was once thought. Actually, it's been linked to diabetes, asthma, neurological problems and birth abnormalities, as well as being a suspected carcinogen. The trouble was that DDT was stupidly stable. It just didn't break down. As a result, it accumulated in the world's ecosystems, ravaging coastal invertebrate wildlife and steadily building up in the tissues of predatory animals. Disturbingly, DDT was still found in human blood samples as recently as 2005 and is regularly found in food samples tested by the FDA.

Sucralose is similarly stable (albeit less so), and when it does break down, it doesn't break into anything harmful. No other sweetener holds the accolade of being considered "safe" by the Center for Science in the Public Interest. The trouble is that it's low calorie because it isn't really absorbed by the human body. Most will pass straight through you, with only around 4 - 12% actually being metabolised. Indeed, the Swedish Environmental Protection Agency has cautioned about potential rising levels of the stuff in wastewater. Water treatment plants were shown to have little effect on sucralose, with it being present at moderate levels in effluent water.

Of course, as I said above, not all organochlorides are toxic as such. I'm certainly not saying that Splenda® is evil. Quite the contrary, it's been found by repeated tests to be perfectly safe at the kind of levels involved in daily food consuption***. Though, knowing how chlorocarbons have caused environmental issues in the past, I can't help but hope that the scientists involved in manufacturing the stuff know exactly what they're unleashing on the world.

Personally, I'll be sticking to natural sugars. If nothing else, in my humble opinion, they taste nicer.



*On that note, I don't take too well to being herded either.

**I say 'still available'. Several have been taken off the market after being found to break down into toxic fragments. Which is still better than in ancient times -- the less said about sugar of lead the better!

***It's worth noting though, that not much data exists regarding higher levels of sucralose, and those data that do exist don't look good. But then, almost anything is toxic at a sufficiently high dosage.

Yep. That's right. An anion made up of uranium and oxygen is called a uranate. UO₂^2-, UO₃^2- and UO₄^2- are all types of uranate. Wait. It gets funnier. If you were to react one of these things in the lab, say, to attach it to a benzene molecule? Technically, that reaction would be called uranation.

I know, I know. It's puerile humour. But puerile humour is just so easy to do when talking about uranates! In fact, uranates tend to be yellow in colour. Conentrated uranium oxide (an intermediate stage in uranium ore processing) is called yellowcake. Despite its unfortunate sounding name, yellowcake is used to prepare fuel for the nuclear reactors in power stations. It's produced by all countries that mine uranium.

As for the uranium ore itself, the most common is known as uranite (mostly uranium dioxide, UO₂). Interestingly, due to uranium decaying by emitting alpha particles, the immensely rare element technetium can be found in uranite ores. Uranites were also the first place helium was discovered on Earth (seeing as an α particle is just the nucleus of a helium atom). See? I can be puerile and interesting at the same time!

Actually, perhaps the most common uranate is a complex ion called diuranate, U₂O₇^2-. The smaller uranate ions tend to lump together into these larger ions. And yes, they're all yellow too. I wonder, if Martin Heinrich Klaproth realised the unfortunate humour in all of this in 1789 when he discovered the element, he might not have named it after the planet William Herschel had discovered 8 years earlier. A planet which, alas, provides similar humour for astronomers. Us scientists can be a childish bunch sometimes.

Perhaps the coolest use of diuranate though, is as an additive to glassware. Uranium glass (sometimes known as vaseline glass) has a striking lime green hue to it. Ultraviolet will also make uranium glass fluoresce. While the uranium is still emitting α particles, the mere 1-2% of uranium found in most modern uranium glass makes it essentially harmless -- barely above everyday background radiation. A sensitive geiger counter will pick it up. A more meagre one probably wouldn't register the difference. Any stray α particles which do actually escape from the glass aren't even capable of penetrating human skin. Uranium used to be used in glassware quite widely before the advent of nuclear technology. These days, it's rather fallen out of favour, though you can still find it in antiques and certain novelty items. And marbles.

Incidentally, those marbles are available for sale on a site called United Nuclear. If not for the fact that there's almost certainly some restriction on importing uranium, I'd be quite tempted...

Because I'm a geek. Obviously.

It seems like NASA's LRO mission has inspired an awful lot of renewed interest in the Moon, recently. Bloggers, twitterers and forum runners alike have all been posting a lot of Moon-related things lately, and personally, I think it's brilliant. Perhaps if enough people are interested, the governments might finally give such ventures more consideration. Hey, I can always hope, right?

One thing a lot of people seem to get hung up over is the question of water, and in honesty, it's a fair thing to be concerned about. Any permanent lunar habitat would obviously need a suitable quantity of water for people to live there. Most evidence points to the fact that the lunar surface is completely devoid of water. This does pose something of a problem.

On the other hand, a lot of people still believe it's open to debate. The idea of ice in deep craters is not a new one. It's a simple concept. Normally with the sunlight that hits the lunar surface, water would either boil off rapidly into space, or be split into hydrogen and oxygen. Which would then boil off into space. The moon isn't a friendly place for small molecules to live. On the other hand, if some craters near the Moon's poles are deep enough, they may never see sunlight. If this is the case, then they may actually remain cold enough to retain some amount of ice indefinitely. After all, the Moon's surface is a chilly -165°C in the shade. That should be just fine, right?

Personally, I'm skeptical of the idea. Even at such a low temperature, a few million years worth of vacuum ablation would surely remove most of not all of the water. Entropy would see to that. Through good old fashioned thermodynamics, I'd expect any reserves of ice, even in deep craters, to have been lost to space by now, molecule by molecule. Though I'll freely admit that there might be something I'm not considering there (thermodynamics and I never really got on that well). All the same, it seems... implausible.

Below the surface? Well, that's another matter. It seems like there might have been water in the Moon once upon a time. How long any pockets of ice could persist are anyone's guess though, especially given the daytime temperature of 110°C. If there are subsurface pockets of water, they'd probably exist as veins, the way minerals do on Earth. Future lunar colonies might need to actually mine for water.

There seems to be one big thing though, that no one's really talked about. Why don't we just make the water once we're there? All the right ingredients are in place (well, hydrogen and oxygen). The trouble is the oxygen is chemically bound up in minerals, and most hydrogen on the moon is from the solar wind. The fact remins though -- If someone could devise a feasible way of isolating the right chemicals and splitting them apart, they could be reacted together to produce water. How easy that would be is a good question. For fairly obvious reasons, no one's ever had to devise a chemical process to make water before.

Perhaps a modified version of the Hall-Héroult process -- the process used on Earth to extract aluminium from alumina (Al₂O₃). That would give the added benefit that it would also produce aluminium, which would be an extremely useful engineering material on the Moon. Alternatively, perhaps it could be as simple as reacting hydrogen gas with nickel oxides.

It may be true that I'm no expert on lunar mineralogy, but as the old saying goes, 'if you can't lead a horse to water...'

I was about to type the phrase "ice is cool", but then I realised what I was saying and decided that even I can't get away with a pun that bad. At least not at the start of a blog post...

As well as being able to keep your drinks cool on a hot day, ice is actually pretty amazing stuff. Even more amazing as it's now been shown that there are 20 different types of ice -- 15 crystalline forms alongside 5 amorphous (non-crystalline) forms. Actually, this is yet another reason why water is pretty special. Very few substances can exist in this many different solid phases.

Different types of ice crystal form according to how water molecules stack together when they solidify. This varies according to the temperature and pressure. For instance, the ice that you'd throw into a glass of iced mocha is known scientifically as ice I_h. It has a hexagonal crystal structure which makes it very low in density. Hence, it floats in your drink because it's actually less dense than liquid water. If you were to take that same ice though, and put it in a pressure chamber, the molecules would start to become unstable in their hexagonal shape. At a certain critical pressure, the molecules would rearrange themselves and the ice would recrystallise into a different structure.

Actually, ice is capable of forming almost every type of crystal structure if you give it the right environment. Monoclinic, orthorhombic, tetragonal, cubic, hexagonal... You name it.

The newest addition to this little crystalline family though, is ice XV, first reported just last week. Thermodynamically stable at temperatures below around 130K (-143°C) and pressures of 0.8-1.5 gigapascals, it's unlikely to be too prolific in nature. There is, however, always a chance that exotic forms of ice like this might be able to exist inside the kind of icy moons and dwarf planets that litter the outer solar system.

It was also previously predicted that ice XV would be ferroelectric -- the kind of conductive material with potential uses in electronics and suchlike. As it happens, the predictions were wrong. Very wrong in fact. Ice XV is actually antiferroelectric!

It's quite surprising when you think about it, that we're still discovering things about something as fundamental on planet Earth as ice!

Source: arXiv Blog, arXiv.

I must confess...

I'm a little nervous about talking to a group of astronomers about molecular spectroscopy. Especially as, frankly, I have utterly no idea how much chemistry everyone actually knows!

Ah, the perils of being interdisciplinary.