Tag Archives: oxygen

Chemical connections: dexamethasone, hydroxychloroquine and rheumatoid arthritis

Posted on by

The chemical structure of dexamethasone (image from Wikimedia Commons)

It’s been widely reported today that a “cheap and widely-available” steroid treatment has been shown to be effective in patients suffering the most severe COVID-19 symptoms, significantly reducing the risk of death for both patients on ventilators and those on oxygen treatment.

Most of the reports have understandably focused on the medical aspects, but this is a chemistry blog (mostly) so *cracks chemistry knuckles* what is dexamethasone, exactly?

Its story starts a little over 60 years ago when, in 1958, a paper was published on “clinical observations with 16a-methyl corticosteroid compounds”. Bear with me, I shall explain. Firstly, corticosteroids are hormones which are naturally produced in our bodies. They do all sorts of nifty, useful things like regulate our immune response, reduce inflammation and help us to get energy from carbohydrates. Two of the most familiar names are probably cortisol and cortisone—both of which are released in response to stress.

The discovery of corticosteroids was an important one. So important, in fact, that a few years earlier, in 1950, Tadeusz Reichstein, Edward Calvin Kendall and Philip Showalter Hench had been awarded a Nobel Prize in Physiology and Medicine for “discoveries relating to the hormones of the adrenal cortex”.

The adrenal glands are two small glands found above the kidneys. The outermost part of these glands is called the adrenal cortex (“cortex” from the Latin for (tree) bark and meaning, literally, an outer layer). In the mid-1930s Kendall and Reichstein managed to isolate several hormones produced by these glands. They then made preparations which, with input from Hench, were used in the 1940s to treat a number of conditions, including rheumatoid arthritis.

This was hugely significant at the time, because until this point the treatments for this painful, debilitating condition were pretty limited. Aspirin was known, of course, but wasn’t particularly effective and long-term use had potentially dangerous side effects. Injectable gold compounds (literally chemical compounds containing Au atoms/ions) had also been tried, but those treatments were slow to work, if they worked at all, and were expensive. The anti-malarial drug, hydroxychloroquine (which has also been in the news quite a lot), had been tried as a “remittive agent”—meaning it could occasionally produce remission—but it wasn’t guaranteed.

Rheumatoid arthritis causes warm, swollen, and painful joints (image from Wikimedia Commons)

Corticosteroids were a game-changer. When Hench and Kendall treated patients with what they called, at the time, “compound E” (cortisone) there was a rapid reduction in joint inflammation. It still caused side effects, and it didn’t prevent joint damage, but it did consistently provide relief from painful symptoms.

Fast-forward to the 1958 paper I mentioned earlier, and scientists had discovered that a little bit of fiddling with the molecular structure of steroid molecules caused them to have different effects in the body. The particular chemical path we’re following here started with prednisolone, which had turned out to be a useful treatment for a number of inflammatory conditions. However, placing a methyl group (—CH3) on the 16th carbon—which is, if you have a look at the diagram below, the one on the pentagon-shaped ring, roughly in the middle—changed things.

The steroid “nucleus”: each number represents a carbon atom (image from Wikimedia Commons)

In 1957, four different molecules with methyl groups on that 16th carbon were made available for clinical trial. One of them was 16a-methyl 9a-fluoroprednisolone, more handily known as dexamethasone.

(Quick aside to explain that on the diagram of dexamethasone at the start of this post, the methyl group on the 16th carbon is represented by a dashed wedge-shape. It’s a 2D diagram of a 3D molecule, and the dashed wedge tells us that the methyl group is pointing away from us, through the paper, or rather, screen. This matters because molecules like this have mirror image forms which usually have very different effects in the body—so it’s important to get the right one.)

Dexamethasone is on the WHO Model List of Essential Medicines

It turned out that dexamethasone had a much stronger anti-inflammatory action than plain prednisolone, and it was also more effective the other molecules being tested. It caused a bigger reduction in symptoms, at lower doses. A win all round. It did still have side effects—weight gain, skin problems and digestive issues—but these were no worse than other steroids, and better than some. In fact, salt and water retention were less with dexamethasone, which meant less bloating. It also seemed to have less of an effect on carbohydrate metabolism, making it potentially safer for patients with diabetes.

Skipping forward to 2020, and dexamethasone is routinely used to treat rheumatoid arthritis, as well as skin diseases, asthma, COPD and various other conditions. It is on the WHO Model List of Essential Medicines—a list of drugs thought to be the most important for taking care of the health needs of the population, based on their effectiveness, safety and relative cost.

In the wake of more and more evidence that COVID-19 disease was leading to autoimmune and autoinflammatory diseases, scientists have been looking at anti-inflammatory drugs to see if any of them might help. The Recovery Trial at the University of Oxford was set up to investigate a few different drugs, including hydroxychloroquine (there it is again) and dexamethasone.

It’s not a miracle cure but, in the most severe cases, dexamethasone—a cheap, 60+ year old drug—might just make all the difference.

And that brings us back to today’s news: in the trial, 2104 patients were given dexamethasone once per day for ten days and compared to 4321 patients who were given standard care. The study, led by Professor Peter Horby and Professor Martin Landray, showed that dexamethasone reduced the risk of dying by one-third in ventilated patients and by one fifth in other patients receiving only oxygen.

It’s not a miracle cure by any means: it doesn’t help patients who don’t (yet) need respiratory support, and it doesn’t work for everyone, but, if you find yourself on a ventilator, there’s a chance this 60+ year-old molecule that was first developed to cure rheumatoid arthritis might, just, save your life. And that’s pretty good news.

EDIT 17th June 2020: Chemistry World published an article pointing out that “the trial results have yet to be released leading some to urge caution when interpreting them” and quoting Ayfer Ali, a specialist in drug repurposing, as saying “we have to wait for the full results to be peer reviewed and remember that it is not a cure for all, just one more tool.

If you’re studying from home, have you got your Pocket Chemist yet? Why not grab one? It’s a hugely useful tool, and by buying one you’ll be supporting this site – it’s win-win!

Like the Chronicle Flask’s Facebook page for regular updates, or follow @chronicleflask on Twitter. Content is © Kat Day 2020. You may share or link to anything here, but you must reference this site if you do. If you enjoy reading my blog, and especially if you’re using information you’ve found here to write a piece for which you will be paid, please consider buying me a coffee through Ko-fi using the button below.
Buy Me a Coffee at ko-fi.com

Want something non-sciency to distract you from, well, everything? Why not check out my fiction blog: the fiction phial.

Practical Pyrotechnics (Happy Birthday, Good Omens!)

Posted on by

The novel, Good Omens, was first published on 10th May 1990.

Today (10th May*) is the thirtieth anniversary of the release of the book Good Omens, which is an old favourite of mine, and one I’ve found science-based excuses to write about before. In honour of the day, I’m going to do it again—but this time I’m going to talk about fire.

Fire plays an important role in both the book and the acclaimed television adaptation. Of course, fire is rather easier to do in a novel, since reading words like “fire” and “flames” are generally quite safe. In TV land, however, it’s a bit trickier. In particular (spoiler alert), at the start of episode five, the bookshop owned by the angel Aziraphale is burning when Crowley arrives and walks in. Crowley, after all, is a demon. From Hell. Fire can’t hurt him.

Except, of course, he’s actually the lovely David Tennant, who is a very much not-fireproof human being. Which poses a few questions: did the film crew really set the bookshop set on fire? Did they really make David Tennant walk into a burning building? How is that done safely? And what did they actually burn?

It turns out that they did, in fact, burn down the bookshop set. According to The Nice and Accurate Good Omens TV Companion, director Douglas Mackinnon “wanted a real fire” and “there were thousands of books, tapestries and beautiful grandfather clocks inside the shop that were real.”

Actual books were harmed in the making of Good Omens (photo used with permission).

Which… argh. Actual books. In flames. I might be a bit traumatised. Give me a moment.

Anyway. The thing is, if you’ve ever set fire to paper you’ll know it’s not very controllable. You can’t just burn books and achieve consistent and, more importantly, safe, flames. The Good Omens TV Companion goes on to explain that the set was rigged with gas lines and flame bars. It doesn’t say what the fuel was, but the probable candidate is propane.

This is where we get to the chemistry. Propane is a hydrocarbon—a molecule made of hydrogen and carbon atoms—and the “prop” part of its name tells us that it contains three carbon atoms. The “ane” part tells us it’s an alkane, and from that, handily, we can work out its formula without having to do anything so mundane as look it up, because the formulas of alkanes follow a rule: CnH2n+2. In other words, take the number of carbons, multiply it by two, add two, and you get the number of hydrogen atoms. This gives us three carbons and eight hydrogens: C3H8.

Propane’s boiling point is -42 oC, meaning it’s a gas at room temperature. You may be familiar with propane canisters which slosh when moved, suggesting liquid, and that’s because the propane is under pressure. The only real difference between a gas and a liquid is the amount of space between the individual particles. In a liquid, the particles are mostly touching one another, while in a gas there are large spaces between them. If you take a gas and squash it into a small volume, so that the particles are forced to touch, it becomes a liquid.

Propane is stored in pressurised canisters (photo used with permission)

But once the propane is allowed to escape from the confines of a pressurised container, at room temperature, its molecules spread out once again, into a gas.

The expansion is BIG. Theoretically, at room temperature, one litre of propane liquid (with a density of 493 g/litre) will expand to occupy roughly 270 litres of space. But, of course, the space it’s expanding into also contains air, so the volume of flammable mixture—approximately 5% propane to 95% air—is actually much higher.

Gases burn faster than either liquids or gases. We know this, of course: it only takes a brief spark to light the gas burner on the cooker hob, for example, but you’d struggle to light a liquid fuel with the same spark (unless it was warmed, and therefore starting to vaporise). The reason is those big gaps between molecules: each molecule in a gas is free, none are “buried” in the middle of a volume of liquid (or solid), so they can all mingle freely with oxygen (needed for combustion) and they all “feel” the heat source and become excited more easily.

Propane is a hydrocarbon with three carbon atoms.

Apart from being a gas at room temperature, propane is also chemically very safe in that it’s non-toxic and non-carcinogenic. It’s also colourless and odourless—although small amounts of additives such as the eggy-smelling ethyl mercaptan (ethanethiol) are sometimes added as a safety precaution, to make leaks more noticeable.

Mechanically there are more hazards. There’s a significant temperature drop when a pressurised liquid expands into a gas. The simplest way to think about this is to think of temperature as the energy of all the particles in a substance divided by its volume. If the volume increases while the number of particles stays the same, the energy is spread out a lot more, so the temperature drops. Potentially, a sudden release of too much gas near a person could severely chill their skin, and even cause frostbite. Plus, of course, although propane isn’t toxic, if it displaces oxygen it could cause asphyxiation, and it’s heavier than air, so it tends to accumulate in the bottom part of a room—precisely where people are trying to do pesky things like breathe.

Yellow flames, and smoke, are a sign of incomplete combustion (photo used with permission).

Then there’s the issue of complete combustion. Generally, when hydrocarbons burn they produce carbon dioxide and water as products, neither of which are too much of a problem for nearby humans (up to a point). However, when there’s not enough oxygen—say, because the fire is inside a building—other products form, in particular carbon monoxide, which is very toxic, and carbon particles, which make a terrible, terrible mess.

I mentioned earlier that a flammable mixture is about 95% air to 5% propane, and this is why. In fact, it’s even more precise than that: for propane to burn cleanly it should be 4.2% propane to 95.8% air. In industry terminology, if there’s not enough propane it produces a “lean” burn, where flames lift from the burner and tend to go out. If there’s more propane (and thus not enough oxygen) it’s called a “rich” burn, which produces large, yellow flames, soot, and the dreaded carbon monoxide.

They did burn the bookshop. But it’s OKAY, it was restored again at the end! (Photo used with permission.)

You might, of course, want a certain amount of yellow flame and smoke, to achieve the right look, but the whole thing needs to be carefully controlled to make sure no one is in danger. It’s all manageable with the use of properly checked, monitored and maintained equipment, but you can imagine that a big effect like the bookshop fire needs a very experienced professional to oversee everything.

For Good Omens, that was Danny Hargreaves (of Real SFX), who’s worked on all kinds of projects from War of the Worlds to Doctor Who. As he says in the Good Omens TV Companion, “everything is under control [but] we took it right to [the] limit.” At one point, he says, he turned off gas lines sooner rather than later and, when director Douglas Mackinnon asked why, had to explain that the roof was about to catch fire.

So, yes, they burned the bookshop set. But it’s all right, everyone. It’s all right. Because (another spoiler) thanks to the powers of Adam Young, everything was restored again afterwards. Phew. All the books were saved. Shh.

*Funnily enough, everyone thought the anniversary was 1st of May. Including the whole Good Omens team. So they made a brilliant lockdown video** to mark the occasion and celebrate. And then it turned out it was actually the 10th. Just an ordinary cock-up, as Crowley would say.

**Which proves the bookshop, with all its books, was fully restored, doesn’t it? Told you.

If you’re studying from home, have you got your Pocket Chemist yet? Why not grab one? It’s a hugely useful tool, and by buying one you’ll be supporting this site – it’s win-win!

Want something non-sciency to distract you? Why not check out my fiction blog: the fiction phial. There are loads of short stories, and even (recently) a couple of poems. Enjoy!

Like the Chronicle Flask’s Facebook page for regular updates, or follow @chronicleflask on Twitter. Content is © Kat Day 2020. You may share or link to anything here, but you must reference this site if you do. If you enjoy reading my blog, please consider buying me a coffee through Ko-fi using the button below.
Buy Me a Coffee at ko-fi.com

What is Water? The Element that Became a Compound

Posted on by

November 2018 marks the 235th anniversary of the day when Antoine Lavoisier proved water to be a compound, rather than an element.

I’m a few days late at the time of writing, but November 12th 2018 was the 235th anniversary of an important discovery. It was the day, in 1783, that Antoine Lavoisier formally declared water to be a compound, not an element.

235 years seems like an awfully long time, probably so long ago that no one knew anything very much. Practically still eye of newt, tongue of bat and leeches for everyone, right? Well, not quite. In fact, there was some nifty science and engineering going on at the time. It was the year that Jean-François Pilâtre de Rozier and François Laurent made the first untethered hot air balloon flight, for example. And chemistry was moving on swiftly: lots of elements had been isolated, including oxygen (1771, by Carl Wilhelm Scheele) and hydrogen (officially by Henry Cavendish in 1766, although others had observed it before he did).

Cavendish had reported that hydrogen produced water when it reacted with oxygen (known then as inflammable air and dephlogisticated air, respectively), and others had carried out similar experiments. However, at the time most chemists favoured phlogiston theory (hence the names) and tried to interpret and explain their results accordingly. Phlogiston theory was the idea that anything which burned contained a fire-like element called phlogiston, which was then “lost” when the substance burned and became “dephlogisticated”.

Cavendish, in particular, explained the fact that inflammable air (hydrogen) left droplets of “dew” behind when it burned in “common air” (the stuff in the room) in terms of phlogiston, by suggesting that water was present in each of the two airs before ignition.

Antoine-Laurent Lavoisier proved that water was a compound. (Line engraving by Louis Jean Desire Delaistre, after a design by Julien Leopold Boilly.)

Lavoisier was very much against phlogiston theory. He carried out experiments in closed vessels with enormous precision, going to great lengths to prove that many substances actually became heavier when they burned and not, as phlogiston theory would have it, lighter. In fact, it’s Lavoisier we have to thank for the names “hydrogen” and “oxygen”. Hydrogen is Greek for “water-former”, whilst oxygen means “acid former”.

When, in June 1783, Lavoisier found out about Cavendish’s experiment he immediately reacted oxygen with hydrogen to produce “water in a very pure state” and prove that the mass of the water which formed was equal to the combined masses of the hydrogen and oxygen he started with.

He then went on to decompose water into oxygen and hydrogen by heating a mixture of water and iron filings. The oxygen that formed combined with the iron to form iron oxide, and he collected the hydrogen gas over mercury. Thanks to his careful measurements, Lavoisier was able to demonstrate that the increased mass of the iron filings plus the mass of the collected gas was, again, equal to the mass of the water he had started with.

Water is a compound of hydrogen and oxygen, with the formula H2O.

There were still arguments, of course (there always are), but phlogiston theory was essentially doomed. Water was a compound, made of two elements, and the process of combustion was nothing more mysterious than elements combining in different ways.

As an aside, Scottish chemist Elizabeth Fulhame deserves a mention at this point. Just a few years after Lavoisier she went on to demonstrate through experiment that many oxidation reactions occur only in the presence of water, but the water is regenerated at the end of the reaction. She is credited today as the chemist who invented the concept of catalysis. (Which is a pretty important concept in chemistry, and yet her name never seems to come up…)

Anyway, proving water’s composition becomes a lot simpler when you have a ready supply of electricity. The first scientist to formally demonstrate this was William Nicholson, in 1800. He discovered that when leads from a battery are placed in water, the water breaks up to form hydrogen and oxygen bubbles, which can be collected separately at the submerged ends of the wires. This is the process we now know as electrolysis.

You can easily carry out the electrolysis of water at home.

In fact, this is a really easy (and safe, I promise!) experiment to do yourself, at home. I did it myself, using an empty TicTac box, two drawing pins, a 9V battery and a bit of baking soda (sodium hydrogencarbonate) dissolved in water – you need this because water on its own is a poor conductor.

The drawing pins are pushed through the bottom of the plastic box, the box is filled with the solution, and then it’s balanced on the terminals of the battery. I’ve used some small test tubes here to collect the gases, but you’ll be able to see the bubbles without them.

Bubbles start to appear immediately. I left mine for about an hour and a half, at which point the test tube on the negative terminal (the cathode) was completely full of gas, which produced a very satisfying squeaky pop when I placed it over a flame.

The positive electrode (the anode) ended up completely covered in what I’m pretty sure is a precipitate of iron hydroxide (the drawing pins presumably being plated steel), which meant that very little oxygen was produced after the first couple of minutes. This is why in proper electrolysis experiments inert graphite or, even better, platinum, electrodes are used. If you do that, you’ll get a 1:2 ratio by volume of oxygen to hydrogen, thus proving water’s formula (H2O) as well.

So there we have it: water is a compound, and not an element. And if you’d like to amuse everyone around the Christmas dinner table, you can prove it with a 9V battery and some drawing pins. Just don’t nick the battery out of your little brother’s favourite toy, okay? (Or, if you do, don’t tell him it was my idea.)

Like the Chronicle Flask’s Facebook page for regular updates, or follow @chronicleflask on Twitter. Content is © Kat Day 2018. You may share or link to anything here, but you must reference this site if you do.

If you enjoy reading my blog, please consider buying me a coffee through Ko-fi using the button below.
Buy Me a Coffee at ko-fi.com

Puking pumpkins: more hydrogen peroxide

Posted on by

It was Halloween yesterday and, unusually for the UK, it fell in school term time. As it turned out, I was teaching chemistry to a group of 12-13 year olds on that day which was too good an opportunity to miss.

Time for the puking pumpkin!

A side note: there’s loads of great chemistry here, and the pumpkin isn’t essential – you could easily do this same experiment during a less pumpkin-prolific month with something else. Puking watermelon, anyone?

Carve a large mouth, draw the eyes and nose with marker pen.

First things first, prepare your pumpkin! Choose a large one – you need room to put a conical flask inside and put the pumpkin’s “lid” securely back in place.

Carve the mouth in the any shape you like, but make it generous. Draw the eyes and nose (and any other decoration) in waterproof marker – unless you want your pumpkin to “puke” out of its nose and eyes as well!

Rest the pumpkin on something wipe-clean (it might leak from the bottom) and put a deep tray in front of it.

To make the “puke” you will need:

  • 35% hydrogen peroxide (corrosive)
  • a stock solution of KI, potassium iodide (low hazard)
  • washing up liquid

The puking pumpkin!

You can also add food colouring or dye, but be aware that the reaction can completely change or even destroy the colours you started with. If colour matters to you, test it first.

Method:

  1. Place about 50 ml (use more if it’s not so fresh) of the hydrogen peroxide into the conical flask, add a few drops of washing up liquid (and dye, if you’re using it).
  2. Add some KI solution and quickly put the pumpkin’s lid back in place.
  3. Enjoy the show!

Check out some video of all this here.

What’s happening? Hydrogen peroxide readily decomposes into oxygen and water, but at room temperature this reaction is slow. KI catalyses the reaction, i.e. speeds it up. (There are other catalysts you could also try if you want to experiment; potassium permanganate for example.) The washing up liquid traps the oxygen gas in foam to produce the “puke”.

The word and symbol equations are:

hydrogen peroxide –> water + oxygen
2H2O2 –> 2H2O + O2

There are several teaching points here:

  • Evidence for chemical change.
  • Compounds vs. elements.
  • Breaking the chemical bonds in a compound to form an element and another compound.
  • Balanced equations / conservation of mass.
  • The idea that when it comes to chemical processes, it’s not just whether a reaction happens that matters, but also how fast it happens…
  • … which of course leads to catalysis. A-level students can look at the relevant equations (see below).

Once the pumpkin has finished puking, demonstrate the test for oxygen gas.

Some health and safety points: the hydrogen peroxide is corrosive so avoid skin contact. Safety goggles are essential, gloves are a Good Idea(™). The reaction is exothermic and steam is produced. A heavy pumpkin lid will almost certainly stay in place but still, stand well back. 

But we’re not done, oh no! What you have at the end of this reaction is essentially a pumpkin full of oxygen gas. Time to crack out the splints and demonstrate/remind your students of the test for oxygen. It’s endlessly fun to put a glowing splint into the pumpkin’s mouth and watch it catch fire, and you’ll be able to do it several times.

And we’re still not done! Once the pumpkin has completely finished “puking”, open it up (carefully) and look inside. Check out that colour! Why is it bluish-black in there?

The inside of the pumpkin is blue-black: iodine is produced which complexes with starch.

It turns out that you also get some iodine produced, and there’s starch in pumpkins. It’s the classic, blue-black starch complex.

Finally, give the outside of the pumpkin a good wipe, take it home, carve out the eyes and nose and pop it outside for the trick or treaters – it’s completely safe to use.

Brace yourselves, more equations coming…

The KI catalyses the reaction because the iodide ions provide an alternative, lower-energy pathway for the decomposition reaction. The iodide reacts with the hydrogen peroxide to form hypoiodite ions (OI–). These react with more hydrogen peroxide to form water, oxygen and more iodide ions – so the iodide is regenerated, and hence is acting as a catalyst.

H2O2 + I– –> H2O + OI–
H2O2 + OI– –> H2O + O2 + I–

The iodine I mentioned comes about because some of the iodide is oxidised to iodine by the oxygen. At this point we have both iodine and iodide ions – these combine to form triiodide, and this forms the familiar blue-black complex.

Phew. That’s enough tricky chemistry for one year. Enjoy your chocolate!

Trick or treat!

 

 

 

Like the Chronicle Flask’s Facebook page for regular updates, or follow @chronicleflask on Twitter. All content is © Kat Day 2017. You may share or link to anything here, but you must reference this site if you do.

All comments are moderated. Abusive comments will be deleted, as will any comments referring to posts on this site which have had comments disabled.

Is oxygen really that good for you?

Posted on by

dove oxygen shampoo officialI don’t find time for huge amounts of television these days, and certainly not adverts. But I recently caught an advert for Dove Oxygen Shampoo out of the corner of my eye, and it brought me up short. Of course, beauty products are full of nonsense generally. Think, for example, of L’Oreal’s famous ingredient, ‘Boswelox’. (A word which, thanks to the wonderful Karl Pilkington, has since acquired a whole new meaning.) A little while ago I wrote a post about a toothpaste that was claiming to contain ‘liquid calcium’ (if it were true, cleaning your teeth would be much more exciting, trust me). It’s just par for the course. Really, is there any point wasting valuable energy continuing to be annoyed by these things?

Well yes, actually. Because this kind of silly hogwash just reinforces the ridiculous ‘science is so terribly hard, oooh aren’t all the complicated words impressive?’ attitude that is so frustratingly prevalent in the world today.

Besides which, picking apart this kind of thing is practically the reason for the existence of this blog. So here goes.

Firstly, a few snippets from Dove’s website:

“Oxygen & Moisture shampoo, conditioner and finishing products are pumped with Oxyfusion Technology, a new generation of moisture. This system moisturises fine, flat hair, giving you hair volume.”

And clicking through a bit further:

“[the shampoo] provides conditioning ingredients fused with oxygen as it instantly dissolves on your hair and breathes life into it.”

Hm.

Let’s start with that last sentence. Firstly: it dissolves on your hair? What does that mean? I’m just going to mention here that the meaning of the word dissolve is taught in year 7 (first year, in old money) science in all secondary schools in this country, and has been for many, many years. So everyone should know it, even the employees of the media company that came up with this tosh. (If you don’t, and you’ve ever muttered anything whatsoever about slipping standards and/or grade inflation, shame on you.)

‘Dissolve’ usually refers to solids. Salt dissolves in water. Sugar dissolves in tea (yes all right, also mostly water). It means that the solid becomes incorporated into the liquid, forming a solution. I haven’t checked, but I’m assuming Dove’s shampoo is not solid, as that would make it rather difficult to get out of the bottle.

Ok, oils and fats dissolve in certain solvents (not water mind you), and they could feasibly be liquid and yet the word still applies. True enough. It’s possible that the original text was ‘dissolves the grease on your hair’ (more or less accurate enough), and some marketing guy said, ‘I like really love it, I really reaaaahhhly do, but can we just lose two words from the middle?’

And yes, I think it’s safe to assume their shampoo mixes with water, because that is quite an important feature of shampoo, but they haven’t said ‘dissolves in the water’, they’ve said ‘dissolves on the hair’, which does sort of give the impression that it’s your hair that’s somehow dissolving the shampoo. Which is just weird.

But misuse of the world dissolve is only a minor irritation. No, my bigger problem is ‘ingredients fused with oxygen’. What the Dove does that mean?

For years and years we’ve been told that oxidants are bad. Or at least, that antioxidants are good (although this hasn’t really been backed up by scientific studies).

Is it difficult to work out that oxygen is an oxidant? It’s the granddaddy of oxidants. It’s the oxidant that all the other oxidants were named after. Oxy/oxi – see?

Chemists have two definitions of oxidation, but they’re broadly equivalent. Oxidation can be thought of as gaining oxygen, or it can be thought of as loss of electrons. Electrons are the negatively-charged particles that surround atoms. I mention them because the phrase ‘free radicals’ often turns up in the same breath as ‘antioxidants’. Free radicals are atoms or molecules which have an unpaired electron. Electrons like to be paired up. They REALLY like to be paired up. When they’re not, they’ll do pretty much anything they can to get paired up. Unpaired electrons are, if you like, the desperate guy at the nightclub at the end of the night. This makes them incredibly reactive, which means they can cause cell damage.

Worse, this happens in a chain reaction – meaning that a single free radical can do an awful lot of harm. So where to antioxidants come in? Well, antioxidants react with free radicals and essentially stop them in their tracks.

oxygen cylinder

Don’t suck on this.

Jolly good. But you see where I’m going here? Oxygen is the complete opposite of this. Yes, we breathe oxygen. It’s quite important stuff. Certainly, if you run out of it you’re in trouble. But it’s far from harmless. The air we breathe is only about twenty-one percent oxygen. Too much oxygen is flat-out dangerous. Breathe air with something like 50% oxygen for any length of time and you risk damaging your lungs, eyes and central nervous system. Really. Hospitals control oxygen use very carefully, and scuba divers who use it have to undergo rigorous training. The fad for oxygen bars has caused real concern in some quarters.

What does ‘ingredients fused with oxygen’ mean? Does it mean Dove have somehow dissolved oxygen in their shampoo? I’m certain that it doesn’t, because this wouldn’t be stable, and it would likely cause your shampoo to ‘go off’ in some way very quickly. Does it mean that their shampoo contains an ingredient that releases oxygen somehow? Hydrogen peroxide famously does this, when it breaks down into oxygen and water. Of course hydrogen peroxide is used to bleach hair, so… probably not (and anyway, again, not stable).

I looked up the ingredients in Dove Oxygen Moisture shampoo (and I’ve reproduced them below). To be honest, looking at the list I’m drawing a blank. My suspicion is that they’re using ‘oxygen’ simply because it’s the latest trendy thing. Oxygen is common enough – water contains one atom of oxygen in every molecule for starters, so they’re safe with the idea that the shampoo contains oxygen in some form – just not elemental oxygen.

But, ok, if I had to pick something… there is an interesting ingredient called ‘guar hydroxypropyltrimonium chloride‘ in there. If that is the one that inspired them, I can see why they went with Oxyfusion Technology – guar hydroxypropyltrimonium chloride hardly trips off the tongue.

Sucrose

Table sugar (sucrose) – perhaps we should wash our hair with this?

620px-Guaran.svg

Guar gum – check your salad dressing. Another conditioning alternative perhaps?

I’ve picked that one out of the list partly because it has ‘hydroxy’ in its name. Now in reality, that just means it contains an -OH group or several. This isn’t anything particularly special, table sugar has eight of ’em. Guar hydroxypropyltrimonium chloride comes from guar gum, which in turn is made from guar beans. Guar gum is a food additive that’s used to thicken foods, and it turns up all over the place (check your salad dressing or ice cream).

Guar hydroxypropyltrimonium chloride has been shown to have conditioning properties, which explains its inclusion in shampoo (this is my other reason for picking it out). It probably does leave your hair feeling nice and soft. And it does have several -OH groups, so it arguably sort of works with the ‘conditioning ingredients fused with oxygen’ claim. In the sense that it has oxygen atoms chemically bonded to it. As does, you know, water.

There’s no way that it releases oxygen though. Now in fairness to Dove, that claim isn’t actually made explicitly anywhere, although the lovely bubbly imagery does its damnedest to imply it.

Bad-Hair-Day

Bad hair day?

And here’s the thing: even if you could, would you want to routinely use a product that releases oxygen directly onto your skin or hair? Given that oxygen is an oxidising agent, and is likely to cause cell damage in high concentrations? Just bear in mind what happens to hair that’s exposed to too much hydrogen peroxide.

And don’t even get me started on the dozens and dozens of moisturisers that claim to do the same. Really? Straight into your skin? There are even some products that claim to do both at once, which frankly is jolly clever. In the Doctor Who sense of clever. I.e. fictional.

But what I want to know is this: after years of anti-oxidant this, and anti-oxidant that, how have we managed to go in exactly the opposite direction without consumers saying ‘er, hang on a minute, surely this has to be a load of old boswelox?’

—

Ingredients in Dove Oxygen & Moisture Shampoo:
Aqua, Sodium Laureth Sulfate, Sodium Chloride, Cocamidopropyl Betaine, Glycerin, Citric Acid, Dimethiconol, Disodium EDTA, Guar Hydroxypropyltrimonium Chloride, Laureth-23, Parfum, PPG-12, TEA-Dodecylbenzenesulfonate, TES-Sulfate, DMDM Hydantoin, Sodium Benzoate, Amyl Cinnamal, Benzyl Alcohol, Benzyl Salicylate, Butylphenyl Methylpropional, Hexyl Cinnamal, Limonene, Linalool, CI 17200, CI 42090.

Why weigh atoms that way?

Posted on by

A couple of days ago I was listening to the latest Radiolab podcast. If you’ve never listened to one of these, you really should. They are beautifully produced and, without fail, utterly fascinating. Over the last year or so I’ve learned about a possible cure for a disease with a 100% mortality rate, an apocryphal Russian story about horses frozen into a block of ice, and a new theory for the end of the dinosaurs where, if I understood it correctly, they were essentially grilled to death. Episodes of Radiolab always feel like a thoroughly good use of brain-time.

Anyway, if you’re still with me and haven’t dashed off to immediately download some of these little gems, the most recent episode is about weights and measures and how we’ve standardised them over the years. In particular the kilogram, which is the last physical standard in use, although possibly not for long (listen to the podcast).

shutterstock_27019597_Alhovik_mod_2

So what are the scales made of…?

This got me thinking about atoms and, in particular, how we decide their mass. This matters you see, because the mass of atoms tells us chemists how much stuff to use. If I want a saline solution with a particular concentration, all I need do is look up the numbers on the periodic table, weigh out the appropriate amount of salt and dilute it with the appropriate amount of water. And if you’re a patient who needs a saline drip, you’d better hope I did it correctly.

Anyway, if you remember your periodic table (which of course you do, but just in case, here’s a picture) all the elements come with two numbers.

The Periodic Table

One of these numbers is the atomic number, which is the number of protons in the nucleus of each atom of the element. Conveniently, nature has managed to produce an atomic nucleus for each number between 1 and, at last count, 118 and if you ‘read’ the periodic table from left to right, top to bottom, you’ll see the numbers go up one at a time.

The other number, relative atomic mass, is a bit less tidy. It still goes up as you go along the periodic table, but in less regular jumps of roughly between one and three.  Without going into lots of detail, relative atomic mass is standardised against 1⁄12 the mass of carbon-12. Which begs the question, why? The more mathematically aware will have clocked that 1⁄12 of 12 is, well, 1. So why don’t we compare all the elements to hydrogen, which actually has a mass of 1? Or if that’s infeasible for some reason why not, I don’t know, choose 1⁄9 of beryllium-9, or 1⁄28 of silicon-28?

Well actually, almost exactly 200 years ago now, atomic mass (called atomic weight, at the time) was originally compared to hydrogen, and it was thought that all elements would have masses which were exact multiples of hydrogen’s.

The problem with this was that as measuring techniques became more sophisticated it became clear that some elements were inconveniently failing to follow the rule. In fact, some were downright contrary, like chlorine which appeared to have a mass which wasn’t even a whole number.

This was, at least partially, sorted out in 1932 when James Chadwick proved the existence of neutrons. The existence of isotopes had already been suggested, but this finally cleared up what the pesky things actually were. It turns out some atoms are fatter than others, having one or two more uncharged particles in their nuclei. This doesn’t change what atom they are – they still have the same number of protons – but it does make them a bit heavier. Take a sample of pure chlorine, for example, and you find that roughly three quarters of the atoms in it have a mass of 35, whereas the other quarter have a mass of 37. These are the isotopes of chlorine: imaginatively named chlorine-35 and chlorine-37. Work out the weighted average of the two and you get 35.5, which is the number you see on periodic tables.

In the mid-20th century something of a minor squabble between chemists and physicists broke out (chemists and physicists often squabble: they’re a bit like the English and the French: they like to visit each other but only so that they can moan about how annoying the other lot are and how badly they do everything). By this time had been a switch from using hydrogen (the lightest element) to oxygen as the standard to which other elemental masses were compared. This was mainly for the convenience of chemical analysis: oxygen combines with a lot of things to make straightforward oxides, whereas hydrides are less common and trickier to work with. Plus, large quantities of hydrogen gas are a bit (in the sense of an elephant being a bit heavy, or cyanide being a bit poisonous) of an explosion risk. Oxygen causes other things to burn jolly nicely, but isn’t actually flammable itself. If you can manage to keep it away from other flammable stuff it’s a far safer option.

The problem was that chemists were using a mass scale based on assigning the number 16 to a natural mixture of oxygen (which contains mostly oxygen-16, with little bits of oxygen-17 and oxygen-18). Physicists, on the other hand, had instead assigned the number 16 to the isotope oxygen-16, which they had isolated using the technique of mass spectrometry.

Josef Mattauch

Physicist Josef Mattauch

You may think the physicists’ method sounds more logical, but the chemists’ reasoning was that in naturally-occurring compounds there would be a mixture of isotopes, so it made sense to use a number based on that mixture since you never actually encounter one atom on its own. Either way, the result was differences in the numbers, admittedly some way down the decimal places, but none the less a difference. Of course it was possible to convert between the two, but at the time scientists were fiddling with such tricksy things as nuclear energy and, of course, bombs. Even a tiny discrepancy in the nth decimal place was potentially catastrophic. Something had to be done.

Edward_Wichers

Chemist Edward Wichers

In 1961 a compromise was agreed, thanks largely to the combined efforts of the physicist Josef Mattauch and the chemist Edward Wichers, who set about persuading their respective groups to be reasonable and play nicely with each other.

The result was that carbon-12 was assigned a mass of exactly 12 and the relative atomic mass scale became based on that. The choice of carbon was, to an extent, somewhat arbitrary. It suited the physicists, who were already using carbon as a standard for mass spectrometry. It fell in between the two previous values (1 for hydrogen and 16 for oxygen), which meant it wouldn’t throw every existing piece of work out by too much. In particular, chemists weren’t keen on switching to the physicists’ method of 1⁄16 of the oxygen-16 isotope, because it would change their numbers quite significantly. Switching to 1⁄12 of carbon-12 meant, surprisingly, a smaller change. Carbon is also, of course, a naturally abundant element and it was easy to get samples of pure carbon.

And that, as they say, is that. The carbon-12 scale is still used today, over 50 years later, and it’s not going anywhere. Hydrogen is officially 1⁄12 the mass of carbon-12, and we use carbon-12 because, basically, it was the only option the chemists and physicists would agree on. Hey, it’s as good a reason as any.

Baffling gases

Posted on by

The benefits of nitrogen tyre inflationToday I had to pay a scary amount of money for new car tyres. And, in yet more evidence that chemistry permeates everywhere, I found this amazing sign in the garage.

I studied it at some length. The diagram on the right particularly fascinated me. They’ve helpfully included a key, which seems to suggest that the peculiarly square-shaped ‘tyre’, labelled as being filled with compressed air, only contains particles of nitrogen vs compressed air signoxygen, water (and water vapour, because ‘water’ isn’t a broad enough label apparently). The other, nitrogen-filled, one appears to contain oxygen, water (and water vapour) and nitrogen. As well as some mysterious green and red circling arrows.

Hm.

I can’t quite get my head around it. Someone drew this, and sent it to printers, and presumably it’s been displayed in more than one reception area (I’m deliberately not naming the specific garage, since the staff there were nice and helpful and gave me a good price really, and I’m sure they had nothing whatsoever to do with the sign beyond being told by Head Office to put it on the wall).

Did no one think to check it with, well, anyone? It’s almost as bonkers wrong as the American school sign advertising ‘leteracy night’.

Ok, so I’m not a car mechanic. My experience in that area is limited to occasionally topping up my own screenwash and once watching my Dad change some spark plugs. But I’m pretty sure that compressed air is, well, compressed air. As the sign itself makes clear, air is about 79% nitrogen and 21% oxygen (the numbers vary, but that bit, at least, is more or less right). Therefore, first problem, if you fill a tyre with compressed air you are by definition filling it with nearly 80% nitrogen.

In fact, I don’t think I’d want to drive a car with tyres which had been filled with pure oxygen and a bit of water, as the key suggests. Oxygen is a jolly effective oxidising agent. Tyres may not be the most flammable things in the world, but I reckon there’s a significant chance that your hot wheels would become a little more literal than you might like.

Moving on, in the round-ish tyre diagram there appears to be water, oxygen and nitrogen. Call me naive, but if you tell me you’re filling my tyres with nitrogen I’m going to assume it’s pure nitrogen. Whereas what you have there is (I’m so sad I counted and worked it out) 40% nitrogen, 27% water 33% oxygen. I dunno what that particular mixture is, but it’s not air and it’s definitely not pure nitrogen.

And then there are the captions underneath: “Undesirable components of the air are removed when the tyre is filled with nitrogen”.  Well not according to that diagram, because there’s still water and oxygen in it…

And, “nitrogen is the only inflation medium developed solely for the use of pneumatic tyre inflation”. What about compressed air then? Doesn’t that count as an ‘inflation medium’?

And, “maintains the correct tyre pressure for longer”. Well, actually I’m not sure about this one (there’s an interesting article here expressing some skepticism though). The Formula One website says that they do indeed fill racing tyres with “a special nitrogen-rich air mixture, designed to minimise variations in tyre pressure with temperature. The mixture also retains the pressure longer than normal air would.” The internet tells me Red Bull alone invested over $100 million in 2012, a significant chunk of which would necessarily have to be research and development. If anyone knows about the best stuff to fill tyres with, it’s Formula One.

But, there is a bit of a difference between cars that are designed to routinely get up to 200 mph and tyres that are designed to cope with temperatures comfortably over 100 oC, and your household runaround that averages 35-ish miles per hour. Let’s say I’m not convinced we’re comparing apples with apples here.

I also don’t understand those blue arrows coming out of the squarish tyre (why IS it that shape anyway?) They seem to suggest that water is escaping. But if so, why is it escaping from one tyre and not the other? And how permeable is rubber to water anyway? (Answer: not very, otherwise the welly boot would have been a bit pointless really.)

The one thing I do accept is that water might, possibly, have a small effect on tyre pressure. Water has an annoying habit of changing state at everyday temperatures. Just 1 millilitre of liquid water occupies over 1300 millilitres when it turns into a gas at atmospheric pressure and 25 oC. If it’s hotter (which it probably would be, if it had suddenly turned into gaseous water) it occupies even more space. Of course it’s more complicated than this because tyres aren’t at atmospheric pressure, but the point stands: if there’s water in your tyres the pressure would fluctuate a bit as they warm up. This’ll happen anyway, since in general gases expand as they warm up, but water could make the difference more significant.

But I checked, when compressed air is produced they take out most of the water. Most of it condenses when the air is compressed, and the condensate is simply removed. Then they use an air dryer and a filter as well. So I dispute the idea that there’s a lot of water in standard compressed air in the first place. (It has since been pointed out – see comments – that a lot of garages have their own air compressors, and that although they’re supposed to dry the air they may not do it very effectively, so there could be a fair bit of water in there, although there shouldn’t be…)

Anyway, my musings over gases were interrupted by having to pay the terribly big bill. It did seem like a lot for about 50 kg of rubber, but I’m assured that tyres with the correct depth of ridgy bits are quite important. They told me they filled them up with nitrogen for free.