Oxygen is everywhere. Really, few elements are more abundant in the universe — in fact just two, hydrogen and helium. It combines with most other elements from the periodic table to form an incredibly wide range of compounds which serve an incredibly wide range of purposes. Just looking at the Earth: oxygen-bearing compounds are found in the mantle, crust, oceans and seas, atmosphere and living organisms, not to mention natural and man-made materials such as silica, zeolites, textiles, ceramics and drugs. Oh, and oxygen also takes part in energy production, as well as a variety of processes that range from metabolic to geological.

Oxygen is most frequently encountered in the form of isotope ¹⁶O, much more stable than ¹⁸O and ¹⁷O — this is because ¹⁶O boasts 8 protons and 8 neutrons, a ‘magic number’ in the atomic world that confers special stability. In this month’s ‘in your element’ article (subscription required), Mark Thiemens from the University of California, San Diego, explains how determining the ratio of oxygen isotopes has greatly contributed to our understanding of the evolution of natural processes and life on Earth. For example, the ratio of ¹⁸O to ¹⁶O is different in the atmosphere and in oceans (this is called the Dole effect). This difference arises from the photosynthesis and respiration of land-based or marine organisms, which means it can be used to deduce the evolution of terrestrial and marine activities.

Yet, the role of oxygen in the formation of the solar system remains unclear. Some meteorites that are known to be among the oldest objects in the solar systems have an unexpected oxygen isotopic distribution. Despite progress in the field, described by Thiemens in his article (they involve measurements on solar wind samples!), this distribution still remains unaccounted for, and exactly how the current celestial objects were formed remains unsolved for now.

Anne

Anne Pichon (Associate Editor, Nature Chemistry)

Calcium is one of the most abundant elements on Earth. It plays various roles in many organisms, whether for the contraction of muscle cells, preserving potential differences across membranes, as a co-factor for some enzymes, or a component of bones and shells, to name a few.

Yet, it is surprisingly scarce in the upper atmosphere. Why could that be? Don’t anxiously skip to the end of this post for the answer… this scarcity remains unexplained for now. In this month’s ‘in your element’ article (subscription required) John Plane, Professor of Atmospheric Chemistry at the University of Leeds, ponders on this mystery.

All of the calcium that is present in the upper atmosphere has actually been brought there by interplanetary dust particles entering the Earth’s atmosphere, in a process called ‘meteoric ablation’. The intriguing data is that the concentration of calcium is much lower than expected — about 200 times lower than that of sodium for example, whereas they are present in roughly the same concentrations in the Earth’s crust. Check out the article to find out how scientists measure metal concentrations in the atmosphere.

Could the interplanetary dust particles be depleted in calcium before they even come in contact with our atmosphere? Could it be that more volatile elements (such as sodium) get ablated from the meteorites much more easily than calcium? Or an effect of a peculiar atmospheric reactivity for calcium? Plane explains how some of these reasons are valid, but only to some extent — and so the depletion in calcium has not yet been entirely accounted for.

Anne

Anne Pichon (Associate Editor, Nature Chemistry)

This month in his ‘in your element’ piece (subscription required) Dante Gatteschi from the University of Florence and the European Institute of Molecular Magnetism describes dysprosium in the same way as love was in La Traviata: “croce e delizia” (a curse and a blessing).

Compounds of rare-earth metals are so similar to each other that it was very tricky to separate, isolate, and identify new rare-earth elements. But Paul Emile Lecoq de Boisbaudran persevered, and when he finally isolated element 66 from its oxide through a time-consuming and multi-step separation he also came up with a most suitable name — from the Greek dys, ‘hard’ and prositos, ‘to get at’. Despite much subsequent research, including in Luigi Rolla’s lab in Florence, dysprosium remained hard to isolate in pure form until the 1950s, when ion exchange techniques came along to facilitate things.

Their diffuse 4f orbitals are mainly responsible for the properties of rare-earth elements — in particular, in some cases, compounds can show magnetic anisotropy. This is an intriguing property that continues to impart dysprosium with some exotic applications. An alloy of dysprosium with iron and terbium will, for example, change size in a varying magnetic field.

Read the article to get a first-hand account on how Dante Gatteschi and his group — some 60 years after Florence had seen much research on rare-earth separations — investigated these magnetic properties to find surprising bulk magnet and single-molecule magnet species. No delizia without croce though, because this story does involve quantum mechanical studies to try and understand these electronic and magnetic behaviours…

Anne

Anne Pichon (Associate Editor, Nature Chemistry)

The first thing most people think of when they hear ‘arsenic’ is ‘poison’. In fact, it has played such a crucial part in many a high-profile murder throughout history that it used to be called ‘poudre de succession’ in French (inheritance powder) — mostly by women, according to the French Wikipédia page (!)

This month (subscription required), Katherine Haxton from Keele University — who also blogs at Endless Possibilities and can be found on Twitter @kjhaxton — explains why arsenic is a particularly suitable element to illustrate the notion that chemicals might be good or bad depending on their use. And so, as arsenic was inheritance powder for the French, Victorians across the Channel in Britain used it for more entertaining purposes — such as self-medication, for example to improve breathing and stamina, to freshen the skin, as an aphrodisiac, and perhaps even an anti-eczema cream by Charles Darwin.

Although organoarsenic coumpounds were prepared as early as the 1750s, their structures remained elusive until the mid-nineteenth century when Bunsen, with some help from Berzelius, identified tetra-methyl-di-arsane — a “fuming liquid with a strong garlic odour” (maybe this is why it remained elusive for so long!).

Although this sounds surprising arsenic has been used in medicine throughout history. Cyclic compounds with As–As bonds for example went on to become relatively efficient drugs against syphilis, especially after a bit of optimization to reduce some side effects and improve handling procedures (air sensitive compounds weren’t the easiest to administer). But despite medical uses, arsenic — fairly abundant in nature, and present in living systems — can readily make its way into ground water and poison large populations. Yet it seems that in 19^th century Austria, people could have consumed about 300 mg of arsenic (more than 4-times the fatal dose) without dropping dead. Could organisms get used to arsenic? Kids, don’t try this at home.

This is a topic that has been thoroughly discussed in light of a recent Science paper describing a bacterium that can supposedly grow by using arsenic instead of phosphorus. Since its publication, it has been widely — and passionately — debated in many venues, including press conferences, magazines, the Internet and as follow-up Technical Comments in Science, and was even mentioned on Nature Chemistry’s very own Blogroll.

Saviour or killer? Check out the article for other anecdotes about arsenic’s ambivalence!

Anne

Anne Pichon (Associate Editor, Nature Chemistry)

In this month’s ‘in your element’ article (subscription required), Thomas Rauchfuss from Illinois points out curious trends in sulfur’s chemistry.

As it turns out, sulfur is a little difficult to describe in a concise manner: although in its elemental form it mostly adopts a crown-shaped 8-membered ring structure, it also exists in 7-membered rings (those are the bright yellow ones) and even traces of smaller rings; and it happily converts to a one-dimensional elastomer when heated. Its anions also like to form chains, which can be extended or reduced, and easily catenated, through redox chemistry.

Similarly, its reactivity can be puzzling, and in particular its catalytic activity. Even though sulfur is well-known to poison industrial catalysts, it actually acts as a catalyst in biological systems — among many other roles. Metal sulfide clusters can quickly transfer electrons, a very desirable property for catalytic functions, and are widespread in biology. Take methanogens, for example, those microbes that produce methane under anaerobic conditions, thus contributing to global warming. Although the precise mechanism continues to intrigue chemists, at least one step involves breaking a methyl–sulfur bond. A reverse reaction, catalyzed by nickel, is now also attracting attention. Another example is the microorganisms that also use metalloenzymes with iron–sulfide sites to convert CO₂ to CO or N₂ to NH₃. Check out the article to find out other sulfur roles.

Even its spelling is a little controversial, with both ‘sulfur’ and ‘sulphur’ widespread in the literature — have a look at our editorial (free but you have to be (freely) registered on nature.com) to find out why we’ve adopted ‘sulfur’ (and nope, this time it’s not simply the Oxford English vs American English spelling, the arguments are more etymologic).

Anne

Anne Pichon (Associate Editor, Nature Chemistry)

It was while making sulfuric acid that Jöns Jacob Berzelius — ‘the father of Swedish chemistry’ — noticed a red residue which he first took for tellurium, as Russell Boyd from Dalhousie University notes in this month’s ‘in your element’ article (subscription required). A more meticulous investigation, however, revealed that the residue displayed different properties, resembling those of sulfur. The new element fell into place between tellurium and sulfur in the chalcogen family of the periodic table, and Berzelius named it selenium (after the Greek word for Moon) owing to its similarity with tellurium (named after the Latin word for Earth).

Although often eclipsed by sulfur in textbooks, selenium has a reactivity of its own. I particularly like the fact that its grey allotrope, the most stable form, conducts electricity better in the light than in the dark, and converts electric current from AC to DC — properties which have not gone unnoticed in the fabrication of photovoltaic cells and rectifiers, respectively. Its red tint also went on to account for a worldwide application: selenium dioxide (which adopts a one-dimensional chain structure) imparts vibrant reds and pinks to glass.

It was only much, much later (140 years after its discovery) that the role of selenium in biological systems was identified. It replaces sulfur in some proteins, which in recent years have been shown to help the prevention of cancer, by hindering radical attacks on cells or possibly also by slowing tumour growth. It is introduced in the body by ingestion — dietary recommendations however follow a fine line between too little and too much, both with potentially very serious consequences. While ingesting too little can lead to a weakened immune system or heart problems, selenium poisoning comes with unpleasant skin or breath odour side effects, can affect mental awareness, and can even be life-threatening at high doses. Around 55 micrograms per day sounds just right — I was very surprised to read that it is contained in a single dried Brazil nut!

Have a look at the article to find out more, but I wouldn’t recommend daily Brazil nut fests.

Anne

Anne Pichon (Associate Editor, Nature Chemistry)

As we announced in this post, we’ll be posting here some anecdotes or characteristics of the element featured each month in the ‘in your element’ section of the journal.

In our June issue, David Lindsay from the University of Reading and William Kerr from the University of Strathclyde write about cobalt — an element thought to be named after evil sprites (kobold in German) — check out the article to find out why! [subscription required]

But cobalt later went on to show its good side. It is an essential trace element in the human body, found in a group of co-enzymes called cobalamins. Vitamin B12, a cobalamin, features the only naturally occurring organometallic bond that cobalt engages in: a cobalt–carbon bond. B12 is pretty crucial for life as it plays a role in the formation of blood as well as the function of the brain and nervous sytem. Oh, and according to Wikipedia, it also treats cyanide poisoning — a use that is hopefully less in demand.

Find out a variety of catalytic characteristics of organocobalt complexes from Lindsay and Kerr’s essay — including a serendipitous discovery that has led to the well-known Pauson–Khand reaction. That being said, Pauson apparently refers to it as the Khand reaction, however. I wonder how Khand calls it?

Anne

Anne Pichon (Associate Editor, Nature Chemistry)

As you might already be aware, each month, someone writes a page in the journal about one element. These short pieces are pretty informal, and often include some anecdotes or historical tales about a particular element. As we make our way through the periodic table, I’ve been wanting to share some of these stories with you.

In our May issue, Ralph Puchta from the University of Erlangen-Nürnberg tells us about beryllium.

Did you know that beryllium plays an important role in the nuclear carbon formation in space? Under just the right conditions, two ₂⁴He nuclei (also known as alpha particles) first combine into a ₄⁸Be atom which can then — despite its instability — form one ₆¹²C atom on encountering a third alpha particle.

It is also present in nature in pretty gemstones such as emeralds and aquamarines, which essentially consist of beryl (beryllium aluminium cyclosilicate), with a few transition metal impurities that give them their colour. Beryl is the mineral from which beryllium was first isolated in 1798 and, obviously, named — although ‘glucinium’ had also been proposed at the time, because beryllium salts tasted sweet. ‘Glucinium’ was finally abandoned after nearly 160 years of using the two names.

Semantics aside, beryllium and a lot of its compounds are known to be toxic (so I wouldn’t want to taste exactly how sweet they are myself!), in particular in the form of powders, and should be handled with care. Still, it displays an array of properties that are attractive for applications ranging from radiation windows for X-ray tubes (it doesn’t absorb X-rays much) to aerospace and military usages (it is light, stiff and resists low temperatures). Beryllium could even soon find its way in the processors of quantum computers.

I’ll let you find out more trivia from Puchta’s article [subscription required to read the article]

Oh – and have we mentioned that we’re running a writing competition based on this ‘in your element’ feature?

We look forward to reading your articles!

Anne

Anne Pichon (Associate Editor, Nature Chemistry)

Since we announced our essay competition last month, we’ve had a few questions regarding eligibility and one or two other queries. We’ve replied individually to each query, but just in case there is any confusion out there, this is what we had to say:

1. Are high-school students eligible?

Yes.

If you are currently a student (at high school, or at university studying for an undergraduate or graduate degree, or at an equivalent institution studying for an equivalent qualification) you are eligible.

2. I’ve finished my undergraduate degree but haven’t started a PhD yet — am I eligible?

Yes, as long as it is no more than five years since you completed your undergraduate degree as of 1st August 2011.

You are also eligible if you are no longer a PhD student or a postdoc — as long as it is no more than five years (on 1st August 2011) since you finished your PhD/postdoc.

3. I’m studying something other than science/chemistry, can I enter?

Yes. You don’t have to be a chemistry/science student to enter, and anyone who is no more than 5 years (on 1st August 2011) from their last formal stint of education — from high-school right up through to postdoc — is eligible to enter.

4. I’ve written something that has been published in a magazine/journal already — can I still enter?

Yes. As long as you fulfill the criteria, you can still enter. What we don’t want are essays from professional science writers who make their main living from science writing. If you’re just starting out on that path (and you still fall under the five-year-rule), then you are more than welcome to submit.

5. Are we allowed to include pictures with our submission?

You can if you wish, it would certainly do no harm to include images.

Try not to make your essay reliant on the picture, however, because should your essay be selected as one of the winning ones, we would then need to make sure the figure would be suitable for publication — and that might lead to complications if we can’t use your suggested image and your essay refers to it a lot.

Getting permissions to use images isn’t always quick and/or easy — and that is why for the In Your Element pieces we have published so far, we have typically used generic royalty-free stock images.

We also recommend that you avoid using technical figures or schematics — these are meant to be easy-read type articles.

With all that said, yes, you are free to include pictures, but you may wish to bear in mind the points made above and be aware that we might not be able to use your suggested image.

And don’t forget that, as well as current affairs in industrial or academic research, we are looking for some anecdotes or interesting stories – perhaps about the element’s history, or its reactivity, or an unusual application. There are some examples of In Your Element articles that we have already published that you can use for guidance (see this post), and we’ve announced who our external judges are.

Good luck!

Anne

Anne Pichon (Associate Editor, Nature Chemistry)