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Saltwater pearls occur mostly along the coast of India, in the Persian Gulf, and in the Red Sea, and are produced by oysters. Freshwater pearls occur in lakes and rivers and are produced by mussels. Oysters can produce only one pearl at a time, while mussels can produce many.



The chemical formulas of conchiolin and calcium carbonate are C30H48O11N9 and CaCO3, respectively. Combined, they are an organic/inorganic compound known as nacre, or mother-of-pearl.

Nacre is 95% calcium carbonate. Calcium carbonate exists within pearls as hexagonal plates roughly half a micrometer thick, a crystalline formation known as aragonite. Nacre is iridescent because the thickness of aragonite crystals is close to the wavelength of visible light.

A brittle ceramic on its own, aragonite lays between fibrous layers of conchiolin. Calcium carbonate assumes the form of aragonite rather than calcite, which represents the bulk of sea shells, because concholin is ionic. This chemical-structural interaction generates impressive strength and is distantly similar to the curing of man-made epoxies. Fossil nautiloid shells unearthed in Oklahoma continue to interfere with light.

Concholin is secreted along with other proteins from the mantles of various mollusks. Concholin is chamber-shaped, enclosing and bonding with aragonite. Functionally, nacre is a kind of permanent mucus; the vast majority of this world's natural pearls are thinly-veneered and expelled shell fragments. The sizeable, round hanks of nacre occasionally produced by oysters are the results of years of uniform irritation. An x-rayed pearl looks like a tree's growth rings.



Despite its relative strength, nacre has experienced almost no functional relationship with humanity.

A pearl dissolved in vinegar is allegedly the most expensive meal in human history, worth eight figures in 2014 US Dollars. This, because of an impromptu wager between Cleopatra and Mark Antony to, well, consume the most expensive meal possible. We're told that Mark Antony could not bring himself to drink. To unfurl every subtext of that act is perhaps beyond my ken, but I can tell you Egypt was Rome's client state in those days, and that Cleo herself stood on an inbred, comparatively incompetent royal bloodline.

Pearls' value in the ancient world is hard to overstate. All cultures with coastline access treasured them obsessively. They were inconceivably rare, even near the Persian Gulf's natural oyster beds. The Koran places pearls among the most desirable objects in Paradise; Ancient Greeks, with their thousand folded miles of beach, displayed pearls at weddings, believing them to absorb love. Documentation of humans' obsession with pearls goes back 4000 years.

In the Hindu religion an appropriate wedding gift is a pearl, undrilled, along with its piercing. As with many attractive and fairly uniform substances, pearls symbolize purity. It was Krishna who discovered the first pearl. Pearls bear much emotional baggage indeed.

Pre-colonial America, North and South, enjoyed a relative abundance of freshwater pearls, and as elsewhere, they were ornamentation for weddings in particular. They were more or less gone from the hemisphere by mid-nineteenth century, being one of the first substances siphoned to Europe.



The culturing of pearls seems a shockingly recent development when one considers everything that goes into gardening, foie gras, and beer. The first cultured pearls appeared in Japan at the turn of the 20th century. Spouses Kokichi and Ume Mikimoto, and biologist Tokichi Nishikawa and carpenter Tatsuhei Mise, each wrapped a grain of sand in oyster epithelial membrane and put the assembly into an oyster.

The Nishikawa/Mise pair won the patent first. The Mise-Nishikawa Method continues to be synonymous with the culturing of pearls. Because Mise took the extra step of patenting a grafting needle, the Mikimiotos were unable to use the Mise-Nishikawa method without invalidating their own patents (they had patented the use of epithelial tissue wrapped around an irritant only). In 1916, the spouses seized upon a technicality, patenting a method to produce round pearls, and Mise and Nishikawa fell into obscurity. Kokichi Mikimoto thenceforth enjoyed a reputation as showman and pest to gem firms and governments.

Today's cultured saltwater pearls are nucleated with, oddly enough, shell bands from American mussels. This was Kokichi's discovery, a result of much trial-and-error. The culturing and inclusion of pearls in jewelry still requires much; thousands of pearls must be sorted through for a single necklace, and drilling requires a machinist's precision and finesse.

Oysters are nucleated at three years of age and hang from rafts suspended in the ocean. Time spent hanging, between one and three years, is determined by the size of pearl desired. Harvest is in Winter. Cultured saltwater pearls are mostly mussel shell with a thin layer of nacre. Freshwater pearls, meanwhile, are all nacre (tissue only is used,) and are more easily induced. Round pearls are the obvious preference in any case, but different shapes can be built around different nuclei.

A pearl's color is vulnerable to water temperature, diet, and mollusc breed. Even today, hatcheries cannot control or predict the colors of finished pearls. So-called "black pearls" are usually very dark blue or green; natural actually-black pearls are known to occurr around French Polynesia.

Not sure if a pearl is real? Rub it against your teeth. The aragonite crystals, smooth to your fingertips, will grate against your tooth enamel.





All About Gemstones. "Natural & Artificial Pearls: Composition & Chemistry.", 12/26/2014.

Human Touch of Chemistry. "How do Oysters Make Pearls?", 12/26/2014.

Pearl. "Composition of Pearl.", 12/26/2014.

Fred Ward, "The History of Pearls.", 12/26/2014.

American Pearl. "A Brief History of Pearls.", 12/26/2014.

Pearl Oasis. "Pearl History.", 12/26/2014.

Pearl Guide. "The History of Pearls.", 12/26/2014.

How Products are Made. "Cultured Pearl.", 12/29/2014.

The Stars and Us

Hydrogen is the smallest and lightest of all elements, but it plays a central role in physics, chemistry and biology. 74% of everything in the universe is hydrogen, by mass; it's more like 90% of all atoms, but each atom is very light, consisting only of one proton and one electron. Another 24% of the mass is helium, which is four times as heavy. In the early universe the proportions were even higher, but some of the hydrogen and helium has been used up in stars, making everything else in the universe. The 2% that isn't hydrogen or helium, in other words. Hydrogen is many stars' single biggest source of power, indeed it's the only source that smaller stars are able to use.

Floating Protons

In chemistry, all standard acids contain hydrogen, and the main measure of acidity is pH, which is determined by the concentration of hydrogen ions (H+) in a solution. A hydrogen ion is an atom of hydrogen that's lost its one electron, leaving behind a single proton: it's a subatomic particle that will turn back into a real atom as soon as it manages to attract a new electron. Acids work because the hydrogen ion engages in a two-pronged attack along with the rest of the acid, which is always a negative ion. For every point that pH increases, the number of hydrogen ions goes down by a factor of ten, so something with a pH of 5 has ten times as many hydrogen ions as something with a pH of 6.

In biology, hydrogen ions are widely used for transporting and storing energy, playing a fundamental role in both respiration and photosynthesis. This 'proton pump' mechanism pushes the hydrogen ions from one side of a membrane to the other, and energy is released when they go back again, like a tiny boulder rolling back down a hill. There's more to it than that, but hopefully you get the gist.

The fact that hydrogen is prone to losing electrons and forming positive ions is one of the things that makes it rather odd, as an element. Usually, only metals do that. But hydrogen is an odd case in many ways; in chemistry teaching I often find myself saying '...except hydrogen.'

Organic Chemistry

It's not only in the form of free protons that hydrogen is important, of course. Hydrogen is found in a huge range of compounds, including almost all organic compounds; it's so ubiquitous that some styles of representing organic compounds just leave out the hydrogen altogether. Instead, places where there isn't hydrogen, but there could have been, are marked out. We're left to fill in the gaps ourselves. That's what's meant by 'unsaturated', when people talk about fats and hydrocarbons - there's some hydrogen missing somewhere, so carbon atoms have to form double bonds with each other, making them less stable and frankly, a whole lot messier. Which is why unsaturated fats aren't solid at room temperature: they don't stack so neatly, so the molecules don't stick together as well.

The hydrogen contributes to the energy density of organic compounds, and having it there enables carbon to form stable chains, rings and so on. Left to its own devices, carbon would mostly only form graphite, diamond and fullerenes, which are all perfectly nice chemicals in their own way, but none of them is anywhere near interesting and versatile to give rise to the complex chemistry of life.

Hydrogen Bonding

The most familiar of all hydrogen compounds is of course water (H2O). It's so much a part of everyday life that it's easy to lose sight of what an odd compound it is. It has the extraordinary property of getting bigger and less dense when it freezes, so unlike almost any other substance its solid form floats on top of its liquid form. If not for this anomaly, Earth's water would almost certainly freeze over, and life would never have had a chance to develop.

Even the fact that water is a liquid at all, at the temperatures we're used to, a little surprising. Most small molecules boil at far lower temperatures. The reason water doesn't has to do with the fact hydrogen is not very good at holding on to electrons, whereas oxygen has exceptional powers of electron-grabbing (it's electronegative, in chemist-speak). That makes for a very uneven molecular partnership, with the oxygen taking the lion's share of each electron for itself. So a water molecule consists of one oxygen atom with the negative charge of almost two whole extra electrons, and two near-naked protons hanging onto one side, with their positive charges exposed to the world. Opposite charges attract, so the hydrogens of one water molecule naturally tend to pull in the oxygen of another. This attraction, known as hydrogen bonding, is why water has such a high boiling point. It's also part of the reason why ice floats. Thanks to hydrogen bonding and the way the hydrogen atoms are arranged around the oxygen, for water to freeze the molecules need to line themselves up in a hexagonal honeycomb pattern, with one water molecule at each vertex, leaving a big gap in the middle.

Hydrogen bonding happens whenever you've got a hydrogen atom attached to oxygen or another electronegative atom, like nitrogen, fluorine or chlorine. It's the single strongest kind of intermolecular force, and it not only explains many of the properties of water, but also those of any other molecule where it features. So ethanol, the alcohol found in booze, is a liquid because it has a hydrogen atom attached to an oxygen (we call this a hydroxyl functional group); but it only has one, so it boils at a lower temperature than water. Glycerol, the alcohol that holds fat together, has three separate hydroxyl groups, so it's got a higher boiling point, and it's very viscous. Glucose has six hydroxyl groups, which is why it's solid at room temperature. All of these are soluble in water, and to some extent also in each other, because hydrogen bonds work between molecules of different types. The same kind of electrostatic forces make water the world's greatest solvent, allowing it to overcome the ionic bonds that hold together many crystals.

Hydrogen Economics

Hydrogen gas is of huge economic importance. The use with the most profound consequences for humankind is the Haber-Bosch process, which turns hydrogen and nitrogen into ammonia (NH3). Nitrogen fixation is such a limiting factor in growing food crops that without this one chemical reaction, we'd only be able to feed about half of the humans currently on this planet.

Producing hydrogen is straightforward, but unfortunately not very efficient. Any acid reacting with a metal will produce hydrogen bubbles, which is how they inflated dirigibles like the Hindenburg, but even the cheapest metals are not that cheap. Electrolysis can be used to split water into hydrogen and oxygen, but that takes quite a bit more energy than you'll ever get back out of it. So most industrially produced hydrogen is currently made from fossil fuels. The usual method is to heat methane with steam in the presence of a catalyst, producing carbon monoxide and hydrogen gas.

For decades people have been talking up the possibility of a hydrogen economy, in which fuel cells would take the place of petrol, and cars run on electrical power, emitting only water. If only we had plentiful hydrogen and good ways to store it and transport it, this would be a very appealing prospect. However, hydrogen is only a liquid at temperatures below about -252°C, and storing it in gas form is impractical unless it's under massive pressure. The challenges of maintaining such temperatures and pressures efficiently may yet be overcome, but given the competition from the developing technologies of algal fuels and battery power, each requiring far less infrastructure investment to take off, it seems likely that hydrogen cells will never fulfil their promise.

More on Hydrogen & Water

Sodium acetate (also known as sodium ethanoate) is currently my favourite chemical. That's only partly because it gets bloody cold in my lab sometimes, and it's nice to have a hand warmer at the ready.

One thing I like about it is how mundane it is. When you react vinegar with sodium bicarb (baking soda) in the classic kitchen chemistry 'volcano' demonstration, it's the chemical left behind once all the carbon dioxide has bubbled away. As you might expect from something easily produced from common food ingredients, it's not particularly harmful. It is a mild irritant though, and we want it much more concentrated than you'd normally get in the kitchen, so wash your hands if you get it on your skin.

The really fun thing about sodium acetate is that you can dissolve far more of it in hot water than in cold - but once it cools down it stays dissolved, until it has something to crystallise around. That's because any change of state requires a nucleation point, a place to get started. In this case, the beginnings of a crystal structure provide a scaffold for more particles to crystallise around. Without that, the solution stays supersaturated, containing more of the chemical than it could normally dissolve. Once a seed crystal is introduced, crystals quickly spread out in all directions, like a cloud of smoke made out of needles of ice. You may have seen this in action, in those re-usable hand warmers - sodium acetate is the main chemical used in them. It only takes a few seconds for the crystals to silently spread through a whole hand warmer or flask, radiating a pleasing warmth as they go.

If you have seen reusable hand warmers in action, you may have wondered where that heat comes from. The answer is that when crystals form, particles are attracted together - electrostatic forces pull them into place, and that gives them energy, which becomes heat. A shoe will drop to the the floor with a thump for essentially the same reason: forces cause things to move, and that movement energy has to go somewhere. So crystallisation is almost always exothermic. The energy it releases is called the latent heat of fusion. By the same token it takes energy to break a crystal's bonds, whether by melting a solid or dissolving it, because that means tearing apart mutually attracted particles. With sodium acetate, this isn't too hard - you can re-dissolve the crystals by just heating the container in boiling water or on a Bunsen burner for a couple of minutes. Anything that heats it up past 58°C for long enough should do the trick.

There are at least two very satisfying ways of demonstrating sodium acetate's power of rapid crystallisation. One way is to introduce a seed crystal to the container, which is what commercial hand warmers do when you pop the little metal thing inside. In the lab I've found that shaking the flask usually triggers the process, thanks to imperfections in the bung, but it's probably more fun to start it off by sprinkling magic dust on top. The other way to do it is to pour the supersaturated solution out onto a seed crystal, creating an instant stalagmite which just keeps growing as long as you keep pouring.

The chemical has various other uses besides making hand warmers and great chemistry demonstrations. It has antifungal and antibacterial properties. It's used as a sealant for concrete. Together with acetic acid, it acts as a buffer, helping keep a solution's pH fairly steady when acids or bases are added to it. Perhaps most excitingly, compounding it with acetic acid allows the formation of sodium diacetate crystals, suitable for flavouring salt and vinegar crisps.

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