At first sight bromine seems to be ‘just another halogen’, a helpful counter-anion or leaving group in S_N2 or cross-coupling reactions. Of course this isn’t the whole story, as Matt Rattley — a chemistry student at the University of Oxford and the author of the winning essay on bromine for last year’s contest — points out in his article (subscription required).

Bromine was isolated independently by Carl Jacob Löwig from a mineral water spring, and Antoine Balard from seaweed, in 1825 and 1826. Having identified that he’d obtained a substance between chlorine and iodine, Balard first thought it was idodine chloride before recognizing it as a new element. It seems unclear who exactly from Balard or Gay-Lussac thought of the name brôme but we know it comes from the Greek bromos, stench — a fair description of gaseous bromine. In addition to its rather unpleasant smell, bromine is also toxic — as Rattley puts it, bromine’s orange-brown colour is convenient because “avoid it you should”. In the sunlight, elemental bromine (Br₂) splits into radicals that readily attack other species, including lung tissues.

Brominated compounds have been used throughout history for a variety of purposes, with varying degrees of success — find out in the article how one was dangerous (likely lethal, really) to ancient Egyptians in the seemingly mundane form of a lipstick. More successful applications include that of potassium bromide, which acts on the nervous system, as an efficient epilepsy remedy, an anticonvulsant and a sedative during the late-19^th and 20^th centuries. It still is used in veterinary medicine, but bromide’s chronic toxicity has since put a stop to human uses. Other instances have also exploited toxicity — Rattley mentions the insecticide chlorenapyr, whose rather peculiar structure comprises three different halogens — while in others no particular problems arose. A polybrominated dye for example has been widely used to stain various cell components for imaging purposes.

In light of such diversity, it certainly doesn’t seem unreasonable to think that bromine will continue to feature prominently both in research and practical applications.

Anne

Anne Pichon (Associate Editor, Nature Chemistry)

Our element of the month is carbon. Carbon is so ubiquitous, with its various allotropes and as part of the many, many compounds and living organisms it makes up, that it’s hard to know where to start. Well, why not in New York City? As Simon Friedman from the University of Missouri Kansas City — interviewed here in Reactions — puts it in his article (subscription required) “The organic chemist’s view of carbon can be like the New Yorker’s view of the world, which to them ends at the edge of Manhattan.” And so he goes on to explain.

The first world of carbon is undeniably organic chemistry, with the incredibly varied species — such as drugs, pesticides, dyes — that it endeavours to synthesize. Yet beside the undisputed, elegant role carbon assumes in organic chemistry, it is also a key component of steel. It is true that iron is by itself a useful material, but it is carbon doping that converts it into steel, an altogether much stronger, much more durable material that can be used to build robust structures. Read Friedman’s article to find out how carbon atoms achieve this.

Another world of carbon — also related to materials and their bulk properties — that has become an inherent part of our lives is plastic. It’s hard to fully grasp just how omnipresent plastics are, from invaluable and advanced items (for example, lenses implanted within the eye) to an unfortunate mountain of junk items filling up landfill sites and even covering a vast area of the Pacific Ocean. Some forms of carbon really are forever, or close enough that we must think carefully about whether this is a good thing or not before (over)using them.

Finally, the last thing you can do with your carbon-based molecules is burn them for energy. We’ve been relying on oil, coal and natural gas for energy — yet in terms of usage of carbon this is more than a little upsetting. I particularly like how Friedman expresses this sentiment: “to the organic chemist, simply burning carbon for its energy must surely be akin to burning your books when you are cold, or eating next year’s seeds when you are hungry”.

Anne

Anne Pichon (Associate Editor, Nature Chemistry)

Readers of this blog will be familiar with last year’s essay competition — as Stuart mentioned here a few days ago, the winning  essay on copper, written by Tiberiu Moga, appears in this month’s ‘in your element’ feature.

Copper has been part of our lives pretty much for ever — the Copper Age started around 5,000 BC (give or take a couple of thousand years depending on whether you count the Copper–Stone Age or not), made its way into epic poetry (read the article to find out how it features in the Kalevala) and copper-based materials are still virtually everywhere, from the humble penny to electrical wiring. So what exactly does copper do, apart from giving her copper(II) carbonate-green colour to the Statue of Liberty? Scientifically speaking, lots of things.

Moga is a Medical Doctor student at the University of Toronto, and previously studied both chemistry and biology at Dartmouth College — he is thus particularly interested in copper’s biological functions and catalytic role in the synthesis of new medicines. He identifies three processes that cover most of its abilities: Lewis-acid catalysis, single-electron-transfer processes, and two-electron-transfer reactions.

One of the best-known reactions involving copper as a Lewis acid is the popular ‘click’ azide–alkyne cycloaddition that connects the two groups to form an azole ring. This fast, reliable reaction is generally easy to carry out and makes for a highly efficient step in a wide variety of processes including, for example, natural product total syntheses.

Single-electron-transfer processes where copper adopts either a Cu⁺ or Cu²⁺ form are widespread in biosystems. Cellular respiration in organisms, for example, relies on a succession of these steps carried out by copper-containing enzymes to oxidize glucose, and extract its energy. Two-electron transfer reactions are also common — they go through a slightly more complex mechanism involving a halide ion.

Of course, this is by no means an exhaustive list. Copper is looking increasingly like a good alternative to palladium catalysts, and it’s also a useful building block — remember the copper nanotubes?

As it turns out, we’re still very much in a copper age, and it looks all set for the duration.

Anne

Anne Pichon (Associate Editor, Nature Chemistry)

This month’s ‘in your element’ article (subscription required) is also a winning entry from last year’s competition. Christine Herman, known on Twitter at @CTHerman, a PhD student at the Department of Chemistry, University of Illinois at Urbana-Champaign, who also likes to write about science — for example she contributes to C&En’s Just Another Electron Pusher — shares why she loves helium.

In 1868, astronomers Jules Janssen and Norman Lockyer — who was about to found a certain Nature journal — both noticed (independently) a bright yellow line in the spectrum of the Sun that could not be accounted for by known elements. The suggestion that this line might come from an element present in the Universe but so far undiscovered on Earth seemed bizarre at first, but was to be later unambiguously backed up. Luigi Palmieri detected this element in 1882 in Mt Vesuvius’ lava, and William Ramsey managed to isolate it in 1895 by treating a sample of the uranium mineral (cleveite) with sulfuric acid, liberating helium that had been produced by the radioactive decay of uranium.

It’s perhaps no wonder that this noble gas wasn’t noticed earlier — it is, after all, colourless, odourless, tasteless, non-toxic, and escapes easily from the Earth’s atmosphere so that its concentration is only about 0.0005% by volume. It does however get trapped under the surface, usually with natural gas, and this is where we get the helium we need.

And need it we do, not just for balloons and squeaky voices at parties. You already know this if you’re, among other things, a paleontologist, a deep-sea diver or an arc welder; read Herman’s article to find out more.

She does make a fair point — helium is cool. So much so that many scientists in many fields (for example physics and medicine but also nuclear energy applications) use it as a cryogen. And if you go down to temperatures below 2 K, helium becomes downright bizarre and very intriguing: it adopts a superfluidic state that has no viscosity but a very high thermal conductivity. It is also enticing to chemists who, undeterred by its inertness, keep trying to combine it with various elements. Some of these — excited dimers rather than actual compounds — went on to find a use in lasers.

And, as if helium wasn’t exciting enough in its own right, antihelium observed last year made for the heaviest anti-particles produced so far. All in all, colourless, odourless, tasteless, non-toxic element 2 is very far from dull.

Anne

Anne Pichon (Associate Editor, Nature Chemistry)

In the midst of the blog relaunch, a trip to China last December (which I plan on telling you about in a future post), and the end-of-year holiday period followed by a start-of-year busy period, I didn’t get the chance to write about our December in your element article. This is the first competition winning essay that we’ve published — I think I did mention last year’s essay competition a couple of times.

Margit Muller – PhD student in pharmacology at the University of Copenhagen – highlights how sodium is far from being as mundane as it may seem. Wise daughters from old fairy tales who tell their royal father they love him as much as sodium chloride (they might have said “salt” in the original version) know this, but let’s take a look at the chemistry arguments.

Since its discovery in 1807 by Sir Humphry Davy – who was on a rather impressive element-discovering spree – sodium has amazed chemists. Reports dating back to the 1850s already describe its spectacular reactivity, including its reaction with water that contributes to entice generations of (mischievous) school kids to chemistry according to some Reactions pieces. Among other applications, it is also what makes for pretty yellow flames in fireworks.

Read the article (subscription required) to find out just how crucial sodium is in biological processes, and how essential it is to maintain a good balance of sodium outside and within the cells. Membrane proteins are in charge of controlling specific sodium channels, which let Na⁺ ions in and out of cell as required and regulate all sorts of processes related to pretty much everything we do, from muscle contraction to neurotransmission. You have been warned, disturbing this sodium influx can have pretty serious consequences! For example, this is just what makes tetrodotoxin from pufferfish (or fugu) — one of the most toxic substances on earth — poisonous…

Anne

Anne Pichon (Associate Editor, Nature Chemistry)

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)