The Chronicles of the Chronicle Flask: 2019

Happy New Year, everyone! Usually, I write this post in December but somehow things have got away from me this year, and I find myself in January. Oops. It’s still early enough in the month to get away with a 2019 round-up, isn’t it? I’m sure it is.

It was a fun year, actually. I wrote several posts with International Year of the Periodic table themes, managed to highlight the tragically-overlooked Elizabeth Fulhame, squeezed in something light-hearted about the U.K.’s weird use of metric and imperial units and discovered the recipe for synthetic poo. Enjoy!

Newland’s early table of the elements

January started with a reminder that 2019 had been officially declared The Year of the Periodic Table, marking 150 years since Dmitri Mendeleev discovered the “Periodic System”. The post included a quick summary of his work, and of course mentioned the last four elements to be officially named: nihonium (113), moscovium (115), tennessine (117) and oganesson (118). Yes, despite what oh-so-many periodic tables still in widespread use suggest (sort it out in 2020, exam boards, please), period 7 is complete, all the elements have been confirmed, and they all have ‘proper’ names.

February featured a post about ruthenium. Its atomic number being not at all significant (there might be a post about rhodium in 2020 😉). Ruthenium and its compounds have lots of uses, including cancer treatments, catalysis, and exposing latent fingerprints in forensic investigations.

March‘s entry was all about a little-known female chemist called Elisabeth Fulhame. She only discovered catalysis. Hardly a significant contribution to the subject. You can’t really blame all those (cough, largely male, cough) chemists for entirely ignoring her work and giving the credit to Berzelius. Ridiculous to even suggest it.

An atom of Mendeleevium, atomic number 101

April summarised the results of the Element Tales Twitter game started by Mark Lorch, in which chemists all over Twitter tried to connect all the elements in one, long chain. It was great fun, and threw up some fascinating element facts and stories. One of my favourites was Mark telling us that when he cleared out his Grandpa’s flat he discovered half a kilogram of sodium metal as well as potassium cyanide and concentrated hydrochloric acid. Fortunately, he managed to stop his family throwing it all down the sink (phew).

May‘s post was written with the help of the lovely Kit Chapman, and was a little trot through the discoveries of five elements: carbon, zinc, helium, francium and tennessine, making the point that elements are never truly discovered by a single person, no matter what the internet (and indeed, books) might tell you.

In June I wrote about something that had been bothering me a while: the concept of describing processes as “chemical” and “physical” changes. It still bothers me. The arguments continue…

In July I came across a linden tree in a local park, and it smelled absolutely delightful. So I wrote about it. Turns out, the flowers contain one of my all-time favourite chemicals (at least in terms of smell): benzaldehyde. As always, natural substances are stuffed full of chemicals, and anyone suggesting otherwise is at best misinformed, at worst outright lying.

Britain loves inches.

In August I wrote about the UK’s unlikely system of units, explaining (for a given value of “explaining”) our weird mishmash of metric and imperial units. As I said to a confused American just the other day, the UK is not on the metric system. The UK occasionally brushes fingers with the metric system, and then immediately denies that it wants anything to do with that sort of thing, thank you very much. This was my favourite post of the year and was in no way inspired by my obsession with the TV adaptation of Good Omens (it was).

In September I returned to one of my favourite targets: quackery. This time it was amber teething necklaces. These are supposed to work (hmm) by releasing succinic acid from the amber beads into the baby’s skin where it… soothes the baby by… some unexplained mechanism. They don’t work and they’re a genuine choking hazard. Don’t waste your money.

October featured a post explaining why refilling plastic bottles might not be quite as simple as you thought. Sure, we all need to cut down on plastic use, but there are good reasons why shops have rules about what you can, and can’t, refill — and they’re not to do with selling more bottles.

November continued the environmental theme with a post was all about some new research into super-slippery coatings that might be applied to all sorts of surfaces, not least ceramic toilet bowls, with the goal of saving some of the water that’s currently used to rinse and clean such surfaces. The best bit about this was that I discovered that synthetic poo is a thing, and that the recipe includes miso. Yummy.

Which brings us to… December, in which I described some simple, minimal-equipment electrolysis experiments that Louise Herbert from STEM Learning had tested out during some teaching training exercises. Got a tic tac box, some drawing pins and a 9V battery? Give it a go!

Well, there we have it. That’s 2019 done and dusted. It’s been fun! I wonder what sort of health scares will turn up for “guilty January”? Won’t be long now…


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Electrolysis Made Easy(ish)

Some STEM Learning trainee teachers, looking very keen!

Back in November last year (was it really that long ago??) I wrote a blog post about water, in which I described a simple at-home version of electrolysis. I didn’t think much of it at the time, beyond the fact that it was oddly exciting to do this experiment—that usually involves power-packs and wires and all sorts of other laboratory stuff—with just a 9V battery, a tic tac box and some drawing pins.

Then, hey, what do you know, someone actually read my ramblings! Not only that, read them and thought: let’s try this. And so it was that Louise Herbert, from STEM Learning (that’s their Twitter, here’s their website), contacted me last month and asked if I’d mind if they used the Chronicle Flask as a source for a STEM learning course on practical work.

Of course not, I said, and please send me some pictures!

And they did, and you can see them scattered through this post. But let’s have a quick look at the chemistry…

Electrolysis is the process of splitting up compounds with electricity. Specifically, ionic compounds: the positively-charged ion in the compound travels to the negative electrode, and the negatively-charged ion moves to the positive electrode.

Water is a covalent compound with the formula H2O, but it does split into ions.

Only… wait a minute… water isn’t ionic, is it? So… why does it work on water? Er. Well. Water does split up into ions, a bit. Not very much under standard conditions, but a bit, so that water does contain very small amounts of OH and H+ ions. (In fact, I can tell you exactly how many H+ ions there are at room temperature, it’s 1×10-7 mol dm-3, and, in an astonishing chemistry plot twist, that 7 you see there is why pure water has a pH of, yep, 7.)

So, in theory you can electrolyse water, because it contains ions. And I’ve more than once waved my hands and left it at that, particularly up to GCSE level (age 16 in the U.K.) because, although it’s a bit of a questionable explanation, (more in a minute), electrolysis is tricky and sometimes there’s something to be said for not pushing students so far that their brains start to dribble out of their ears. (As the saying goes, “all models are wrong, but some are useful.”)

Chemists write half equations to show what the electrons are doing in these sorts of reactions and, in very simple terms, we can imagine that at the positive electrode (also called the anode) the OH ions lose electrons to form oxygen and water, like so:

4OH —> 2H2O + O2 + 4e

And conversely, at the negative electrode (also called the cathode), the H+ ions gain electrons to form hydrogen gas, like so:

2H+ + 2e —> H2

These equations balance in terms of species and charges. They make the point that negative ions move to the anode and positive ions move to the cathode. They match our observation that oxygen and hydrogen gases form. Fine.

Except that the experiment, like this, doesn’t work very well (not with simple equipment, anyway), because pure water is a poor electrical conductor. Yes, popular media holds that a toaster in the bath is certain death due to electrocution, but this is because bathwater isn’t pure water. It’s all the salts in the water, from sweat or bath products or… whatever… that do the conducting.

My original experiment, using water containing a small amount of sodium hydrogen carbonate.

To make the process work, we can throw in a bit of acid (source of H+ ions) or alkali (source of OH ions), which improves the conductivity, and et voilĂ , hydrogen gas forms at the cathode and oxygen gas forms at the anode. Lovely. When I set up my original 9V battery experiment, I added baking soda (sodium hydrogencarbonate), and it worked beautifully.

But now, we start to run into trouble with those equations. Because if you, say, throw an excess of H+ ions into water, they “mop up” most of the available OH ions:

H+ + OH —> H2O

…so where are we going to get 4OH from for the anode half equation? It’s a similar, if slightly less extreme, problem if you add excess alkali: now there’s very little H+.

Um. So. The simple half equations are… a bit of a fib (even, very probably, if you use a pH neutral source of ions such as sodium sulfate, as the STEM Learning team did — see below).

What’s the truth? When there’s plenty of H+ present, what’s almost certainly happening at the anode is water splitting into oxygen and more hydrogen ions:
2H2O —>  + O2 + 4H+ + 4e

while the cathode reaction is the same as before:
2H+ + 2e —> H2

Simple enough, really, but means we use the “negative ions are going to the positive electrode” thing, which is tricky for GCSE students, who haven’t yet encountered standard electrode potentials, to get their heads around, and this is why (I think) textbooks often go with the OH-reacts-at-the-anode explanation.

Likewise, in the presence of excess alkali, the half equations are probably:

Anode: 4OH —> 2H2O + O2 + 4e
Cathode: 2H2O + 2e —> H2 + OH

This time there is plenty of OH, but very little H+, so it’s the cathode half equation that’s different.

Taking a break from equations for a moment, there are some practical issues with this experiment. One is the drawing pins. Chemists usually use graphite or platinum electrodes in electrolysis experiments because they’re inert. But good quality samples of both are also (a) more difficult and more expensive to get hold of and (b) trickier to push through a tic tac box. (There are examples of people doing electrolysis with pencil “leads” online, such as this one — but the graphite in pencils is mixed with other compounds, notably clay, and it’s prone to cracks, so I imagine this works less often and less well than these photos suggest.)

A different version of the experiment…

Drawing pins, on the other hand, are made of metal, and will contain at least one of zinc, copper or iron, all of which could get involved in chemical reactions during the experiment.

When I did mine, I thought I was probably seeing iron(III) hydroxide forming, based, mainly, on the brownish precipitate which looked fairly typical of that compound. One of Louise’s team suggested there might be a zinc displacement reaction occurring, which would make sense if the drawing pins are galvanized. Zinc hydroxide is quite insoluble, so you’d expect a white precipitate. Either way, the formation of a solid around the anode quickly starts to interfere with the production of oxygen gas, so you want to make your observations quickly and you probably won’t collect enough oxygen to carry out a reliable gas test.

In one of their experiments the STEM Learning team added bromothymol blue indicator (Edit: no, they didn’t, oops, see below) to the water and used sodium sulfate as (a pH neutral) source of ions. Bromothymol blue is sensitive to slight pH changes around pH 7: it’s yellow below pH 6 and blue above pH 7.6. If you look closely at the photo you can see that the solution around the anode (on the right in the photo above, I think *squint*) does look slightly yellow-ish green, suggesting a slightly lower pH… but… there’s not much in it. This could make sense. The balanced-for-H+ half equations would suggest that, actually, there’s H+ sloshing around both electrodes (being formed at one, used up at the other), but we’re forming more around the anode, so we’d expect it to have the slightly lower pH.

The blue colour does, unfortunately, look a bit like copper sulfate solution, which might be confusing for students who struggle to keep these experiments straight in their heads at the best of times. One to save for A level classes, perhaps.

(After I published this, Louise clarified that the experiment in the photo is, in fact, copper sulfate. Ooops. Yes, folks, it looks like copper sulfate because it is copper sulfate. But I thought I’d leave the paragraph above for now since it’s still an interesting discussion!)

The other practical issue is that you need a lot of tic tac boxes, which means that someone has to eat a lot of tic tacs. There might be worse problems to have. I daresay “your homework is to eat a box of tic tacs and bring me the empty box” would actually be quite popular.

So, there we are. There’s a lot of potential (haha, sorry) here: you could easily put together multiple class sets of this for a few pounds—the biggest cost is going to be a bulk order of 9V batteries, which you can buy for less than ÂŁ1 each—and it uses small quantities of innocuous chemicals, so it’s pretty safe. Students could even have their own experiment and not have to work in groups of threes or more, battling with dodgy wires and trippy power-packs (we’ve all been there).

Why not give it a try? And if you do, send me photos!


Like the Chronicle Flask’s Facebook page for regular updates, or follow @chronicleflask on Twitter. Content is © Kat Day 2019 (photos courtesy of STEM Learning UK and Louise Herbert). 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.
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Non-stick toilets, synthetic poo and saving the environment

141 billion litres of water are used to flush toilets every day.

Scientists develop slippery toilet coating that stops poo sticking,” shouted newspaper headlines last week, naturally prompting comments about the state of politics, the usual arguments about the ‘right’ way to hang toilet paper rolls, and puns of varying quality.

There was also more than one person asking WHY, given everything going on at the moment, scientists are spending their time on something which seems, well, not terribly urgent. After all, ceramic toilet bowls are already quite slippery. Toilet brushes exist. We have a myriad of toilet cleaning chemicals. Surely there are higher priorities? Attempting to deal with looming environmental disaster, say?

But here’s the thing, from an environmental point of view, flush toilets are quite significant. If you’re fortunate enough to live somewhere they’re ubiquitous it’s easy to take them for granted, but consider this: flushing even a water-efficient toilet uses at least five litres of water (much more for older models, a bit less if you use a ‘half-flush’ function). Often this is perfectly clean water which has been through water treatment, only to be immediately turned back into, effectively, sewage. Now imagine you have something a bit… ahem… sticky to flush. What do you do? You flush the toilet twice. Maybe more. You break out the toilet brush and the bottle of toilet cleaner, and then you probably flush at least one extra time to leave the bowl clean.

Using toilet cleaning chemicals often results in extra flushes.

Consider that the average person uses the toilet about five times and day and multiply up by the population and, even just in the UK, we’re looking at billions of litres of water daily. Globally, it’s estimated that 141 billion litres of fresh water are used daily for toilet flushing, and in some homes it could account for a quarter of indoor wastewater production. That’s a lot of fresh water we’re chucking, quite literally, down the toilet.

It rains a fair bit in the U.K. so, except for the occasional dry summer, Brits aren’t in the habit of worrying too much about water supply. The opposite, if anything. But we need to change our ways. In a speech in March this year, Sir James Bevan, Chief Executive of the Environment Agency, warned that the U.K. could run into serious water supply problems in 25 years due to climate change, population growth and poor water management.

Even putting those warnings to one side, treating water uses energy and resources. Filters are used which have to be cleaned and replaced, chemical coagulants and chlorine (usually in the form of low levels of chlorine dioxide) have to be added. Sometimes ozone dosing is used. The pH of the water needs to be checked and adjusted. All of these chemicals have to be produced before they’re used to treat the some 17 billion litres of water that are delivered to UK homes and businesses every day. And, of course, the whole water treatment process has to be continuously and carefully monitored, which requires equipment and people. None of this comes for free.

So, yes, saving fresh water is important. Plugging leaks and using water-saving appliances is vital. And, given that everyone has to go to the toilet several times a day, making toilets more efficient is potentially a really significant saving. An super non-stick toilet surface could mean less flushing is needed and, probably, fewer cleaning products too — saving chemical contamination.

Fresh water is a valuable resource.

The new super-slippery surface was co-developed by Jing Wang in the Department of Mechanical Engineering at the University of Michigan. It’s called a liquid-entrenched smooth surface (LESS) and is applied in two stages. First, a polymer spray, which dries to form nanoscale hair-like strands. The second spray completely covers these ‘hairs’ with a thin layer of lubricant, forming an incredibly flat, and very slippery, surface. The researchers tested the surface with various liquids and synthetic faecal matter and the difference — as seen in the video on this page — is really quite astonishing.

Hold up a moment, synthetic faecal matter? I’ll bet no one embarking on an engineering degree ever imagines that, one day, they might be carefully considering the make-up of artificial poo. But actually, when you think about it, it’s quite important. Quite aside from safety aspects and the sheer horror of the very idea, you couldn’t use the real thing to test something like this. You need to make sure it has a carefully-controlled consistency, for starters. It’s the most basic principle, isn’t it? If you want to test something, you have to control your variables.

Artificial poo is surprisingly important.

Indeed, there’s even a scale. It’s called the Bristol stool scale, and it goes from “hard” to “entirely liquid”. Synthetic poo is a mixture of yeast, psyllium, peanut oil, miso (proof, if it were needed, that miso really does improve everything), polyethylene glycol, calcium phosphate, cellulose and water. The amount of water is adjusted to match different points on the Bristol scale. Aren’t science and engineering fun?

Anyway. Back to the non-stick technology. This new surface can be applied to all sorts of materials including ceramic and metal, and it repels liquids and ‘viscoelastic solids‘ (stuff that’s stretchy but also resists flow: apart from poo, PVA slime is another example) much more effectively than other types of non-stick surfaces. In fact, the researchers say it’s up to 90% more effective than even the best repellent materials, and they estimate that the amount of water needed to clean a surface treated in this way is 10% that needed for ordinary surfaces. They were also able to show that bacteria don’t stick to LESS-coated materials, meaning that even if untreated water is used to flush a toilet, it remains hygienic without the need for extra chemicals.

The potential to cut 141 billion litres of water by a factor of ten is not to be (I’m sorry) sniffed at. Plus, in some areas, ready supplies of water and the facilities to clean toilets just aren’t available. Using LESS could, potentially, reduce the spread of infection.

By Chemystery22 - Own work, CC BY-SA 3.0, https://commons.wikimedia.org/w/index.php?curid=31161897 A graft copolymer has side chains branching off the main chain — these side chains are the “hairs” described by the researchers.

So what IS this surface treatment made of? This information wasn’t widely reported, but it seems quite important, not least because applications of LESS are estimated to last for about 500 flushes, which suggests that re-application will be needed fairly regularly and, perhaps more worryingly, whatever-it-is is passing into the wastewater supply.

Not surprisingly, there’s a certain amount of vagueness when it comes to its exact make-up, but I did find some details. Firstly, it’s what’s known as a graft polymer, that is, a polymer chain with long side chains attached — these are the “hairs” described by the researchers.

Secondly, the polymer strands are based on polydimethylsiloxane, or PDMS. This may sound terrifying, but it’s really not. PDMS (also known as dimethicone) is a silicone — a compound made up of silicon, oxygen, carbon and hydrogen. These compounds turn up all over the place. They’re used contact lenses, shampoos, and even as food additives. Oh, and condom lubricants. So… pretty harmless. In fact, they’re reported as having no harmful effects or organisms or the environment. The one downside is that PDMS isn’t biodegradable, but it is something that’s absorbed at water treatment facilities already, so nothing new would need to be put in place to deal with it.

The problem of better toilets might be more urgent than you thought.

Finally, the lubricant which is sprayed over the polymer chains in the second stage of the treatment to make the surface “nanoscopically smooth” (that is, flat on a 1 billionth of a metre scale) is plain old silicone oil, which is, again, something with a low environmental impact and generally considered to be very safe.

As always with environmental considerations it’s about choosing the least bad option, and using these coatings would certainly seem to be a far better option than wasting billions of gallons of precious fresh water.

In short, silly headlines aside, it turns out that making toilets better might be quite an important problem. Maybe it’s time to rage against the latrine.


Like the Chronicle Flask’s Facebook page for regular updates, or follow @chronicleflask on Twitter. Content is © Kat Day 2019. 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.
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How are amber teething necklaces supposed to work?

Do amber beads have medicinal properties?

Amber, as anyone that was paying attention during Jurassic Park will tell you, is fossilised resin from trees that lived at least twenty million years ago (although some scientists have speculated it could be older). It takes the form of clear yellow through to dark brown stones, seemingly warm to the touch, smooth and surprisingly hard. It is certainly beautiful. But does it also have medicinal properties? And if it does, are they risk-free?

In 2016 a one year-old boy was found dead at his daycare centre in Florida. The cause of death was a necklace, which had become tangled and tightened to the point that he was unable to breathe.

Why was he wearing a necklace? Surely everyone knows that babies shouldn’t wear jewellery around their necks where it could so easily cause a terrible tragedy like this? No one needs a necklace, after all – it’s purely a decorative thing. Isn’t it?

Yes. Yes, it is. However, this particular type of jewellery was specifically sold for use by babies. Sold as a product that parents should give their children to wear, despite all the advice from medical professionals. Why? Because this jewellery was made from amber, and that’s supposed to help with teething pains.

Teething is a literal pain.

Anyone whose ever had children will tell you that teeth are basically a non-stop, literal pain from about 4 months onward. Even once your child appears to have a full set, you’re not done. The first lot start falling out somewhere around age five, resulting in teeth that can be wobbly for weeks. And then there are larger molars that come through at the back somewhere around age seven. Teenagers often find themselves suffering through braces and, even when all that’s done, there’s the joy of wisdom teeth still to come.

It’s particularly difficult with babies, who can’t tell you what hurts and who probably have inconsistent sleep habits at the best of times. Twenty sharp teeth poking through swollen gums at different times has to be unpleasant. Who could blame any parent for trying, well, pretty much anything to soothe the discomfort?

Enter amber teething necklaces. They’re sold as a “natural” way to soothe teething pain. They look nice, too, which I’m sure is part of their appeal. A chewed plastic teething ring isn’t the sort of thing to keep in baby’s keepsake box, but a pretty necklace, well, I’m sure many parents have imagined getting that out, running their fingers over the beads and having a sentimental moment years in the future.

Amber is fossilised tree resin.

So-called amber teething necklaces are made from “Baltic amber,” that is, amber from the Baltic region: the largest known deposit of amber. It is found in other geographical locations, but it seems that the conditions – and tree species – were just right in the Baltic region to produce large deposits.

Chemically, it’s also known as succinite, and its structure is complicated. It’s what chemists would call a supramolecule: a complex of two or more (often large) molecules that aren’t covalently bonded. There are cross-links within its structure, which make it much denser than you might imagine something that started as tree resin to be. Baltic amber, in particular, also contains something else: between 3-8% succinic acid.

Succinic acid is a dicarboxylic acid.

Succinic acid is a much simpler molecule with the IUPAC name of butanedioic acid. It contains two carboxylic acid groups, a group of atoms we’re all familiar with whether we realise it or not – because we’ve all met vinegar, which contains the carboxylic acid also known as ethanoic acid. If you imagine chopping succinic acid right down the middle (and adding a few extra hydrogen atoms), you’d end up with two ethanoic acid molecules.

Succinic acid (the name comes from the Latin, succinum, meaning amber) is produced naturally in the body where it is (or, rather, succinate ions are) an important intermediate in lots of chemical reactions. Exposure-wise it’s generally considered pretty safe at low levels and it’s a permitted food additive, used as an acidity regulator. In European countries, you might see it on labels listed as E363. It also turns up in a number of pharmaceutical products, where it’s used as an excipient – something that helps to stabilise or enhance the action of the main active ingredient. Often, again, it’s there to regulate acidity.

Basically, it’s mostly harmless. And therefore, an ideal candidate for the alternative medicine crowd, who make a number of claims about its properties. I found one site claiming that it could “improve cellular respiration” which… well, if you’ve got problem with cellular respiration, you’re less in need of succinic acid and more in need of a coffin. Supposedly it also relives stress and prevents colds, because doesn’t everything? And, of course, it allegedly relieves teething pains in babies, either thanks to its general soothing effect or because it’s supposed to reduce inflammation, or both.

Purporters claim succinic acid is absorbed through the skin.

The reasoning is usually presented like this: succinic acid is released from the amber when the baby wears the necklace or bracelet and is absorbed through the baby’s skin into their body, where it works its magical, soothing effects.

Now. Hold on, one minute. Whether this is true or not – and getting substances to absorb through skin is far less simple than many people imagine, after all, skin evolved as a barrier – do you really, really, want your baby’s skin exposed to a random quantity of an acidic compound? Succinic acid may be pretty harmless but, as always, the dose makes the poison. Concentrated exposure causes skin and eye irritation. Okay, you might say, it’s unlikely that an amber necklace is going to produce anywhere near the quantities to cause that sort of effect, but if that’s your logic, then how can it also produce enough to pass through skin and have any sort of biological effect on the body?

The answer, perhaps predictably, is that it doesn’t. In a paper published in 2019, a group of scientists actually went to the trouble of powdering Baltic amber beads and dissolving the powder in sulfuric acid to measure how much succinic acid they actually contained. They then compared those results with what happened when undamaged beads from the same batches were submerged in solvents, with the aim of working out how much succinic acid beads might conceivably release into human skin. The answer? They couldn’t measure any. No succinic acid was released into the solvents, at all. None.

Scientists submerged Baltic amber beads in solvents to see how much succinic acid they released.

They concluded that there was “no evidence to suggest that the purported active ingredient succinic acid could be released from the beads into human skin” and also added that they found no evidence to suggest that succinic acid even had anti-inflammatory properties in the first place.

So amber necklaces don’t work to relieve teething pains. They can’t. Of course, there could be a sort of placebo effect – teething pain is very much one of those comes-and-goes things. It’s very easy to make connections that just aren’t there in this kind of situation, and imagine that the baby is more settled because of the necklace, when in fact they might have calmed down over the next few hours anyway. Or maybe they’re just distracted by the pretty beads.

And, fine. If wearing the jewellery was really risk-free, then why not? But as the story at the start of this post proves, it is not. Any kind of string around a baby’s neck can become twisted, interfering with their breathing. Most necklaces claim to have some sort of “emergency release” mechanism so that they come apart when pulled, but this doesn’t always work.

Don’t fall for the marketing.

Ah, goes the argument. But it’s okay, because we only sell bracelets and anklets for babies. They don’t go around the baby’s neck. It’s completely safe!

No. Because I don’t care how carefully you make it: the string or cord could still break (especially if it’s been chewed), leaving loose beads to pose a serious choking hazard. Not to mention get jammed in ears or nostrils. Even if you’re with the baby, watching them, these sorts of accidents can happen frighteningly quickly. Letting a baby sleep with such an item is nothing short of asking for disaster, and no matter how good anyone’s intentions, babies do have a habit of dozing off at odd times. Will you really wake the child up to take off their bracelet? Every time?

In summary, don’t fall for the marketing. Amber necklaces may be pretty, but they’re not suitable for babies. The claims about succinic acid are completely baseless, and the risks are very real.


Like the Chronicle Flask’s Facebook page for regular updates, or follow @chronicleflask on Twitter. Content is © Kat Day 2019. 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.
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A natural remedy that’s full of chemicals?

Blossoms

The summer holidays are here! A time when parents of small children find themselves exploring every park in their local vicinity, quite probably several times (whilst hoping against hope that it doesn’t rain). On just such a quest myself, I recently visited one particular park that was filled with a gorgeous smell.

What was it? A bit of sniffing around quickly identified this tree. Now, I am not a botanist (or even much of a gardener), so I immediately resorted to the rather wonderful Seek app by iNaturalist, which uses some very clever image recognition software to identify plants and animals (disclaimer: accuracy is not guaranteed — don’t eat anything based on this app!)

Seek told me that this was a lime tree, or a linden (genus Tilia). A bit of cross-referencing (thanks Dad!) suggested that it had identified the tree correctly. It’s not an uncommon plant: you’ll probably come across it yourself if you go looking (or smelling).

The name ‘linden’ was more familiar to me. The wood is soft and easily worked, and is used to make musical instruments because it has good acoustic properties. It’s also used to make wooden blinds and other pieces of furniture because it’s lightweight, stable, and holds stains and finishes well.

Linden blossoms can be used to make tea.

But let’s go back to the flowers and their delicious scent. The tree blooms during July and August in the Northern hemisphere. The flowers are sometimes described as mucilaginous — which is a fabulous word meaning, basically, thick and sticky. More specifically: “containing a polysaccharide substance that is extracted as a viscous or gelatinous solution and used in medicines and adhesives.”

Linden flowers are a ‘natural remedy’ with a list of applications in herbal medicine as long as your arm. They contain lots of different substances. One that comes up a lot is farnesol, which is actually a type of alcohol. Of course, it’s nothing like the alcohol we’re familiar with from drinks, which is the much simpler ethanol — but it’s important to remember that ‘alcohol’ actually refers to a class of compounds (which, in simple terms, contain an -OH group like the one in the image here) and not a single substance.

The chemical structure of farnesol

Farnesol turns up in lots of essential oils, such as citronella, rose and lemon grass. It’s used in perfumes to enhance floral scents. But plants don’t make substances just to please humans (well, it’s complicated…). It acts as a pheromone for several insects. Sometimes this doesn’t work out so well for the insects, as it confuses their mating behaviour and effectively acts as a natural pesticide. On the other hand, it actively encourages others: bumblebees release farnesol when they return to the hive to spur other bees into action. It’s the bee equivalent of shouting, ‘oi! Move it you lot, pollen this way!’

Farnesol acts as a pheromone for bumblebees.

Linden flowers also contain one of my all-time favourite chemicals, benzaldehyde. That’s the one that smells of almonds and isn’t a deadly cyanide salt. Its delicious almondy-ness is the reason it’s used as a flavouring and scent, but it’s also a starting material for loads of different chemicals, for example the dye malachite green, which is used to give a green colour to leather, fabric and paper. A form of this dye called ‘brilliant green‘ is mixed with a second, violet, dye to make ‘Bonney’s blue,’ a disinfectant dye used to mark skin for surgeries. Benzaldehyde is also used to make styrene, which is of course used to make the well-known packing material, polystyrene.

And these are just a couple of the substances found in those yummy-smelling flowers. They also contain arabinogalactans, uronic acid, tannins, rutin, hyperoside, quercitrin, isoquercitrin, astragalin and others. In short, a veritable cocktail of different chemicals.

So next time you smell the scent of a lovely flower, just think about all the amazing chemical substances the plant is making. All natural, of course!


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What is Water? The Element that Became a Compound

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 Pilùtre 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.)


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Carbon dioxide: the good, the bad, and the future

Carbon dioxide is a small molecule with the structure O=C=O

Carbon dioxide has been in and out of the news this summer for one reason or another, but why? Is this stuff helpful, or heinous?

It’s certainly a significant part of our history. Let’s take that history to its literal limits and start at the very beginning. To quote the great Terry Pratchett: “In the beginning, there was nothing, which exploded.”

(Probably.) This happened around 13.8 billion years ago. Afterwards, stuff flew around for a while (forgive me, cosmologists). Then, about 4.5 billion years ago, the Earth formed out of debris that had collected around our Sun. Temperatures on this early Earth were extremely hot, there was a lot of volcanic activity, and there might have been some liquid water. The atmosphere was mostly hydrogen and helium.

The early Earth was bashed about by other space stuff, and one big collision almost certainly resulted in the formation of the Moon. A lot of other debris vaporised on impact releasing gases, and substances trapped within the Earth started to escape from its crust. The result was Earth’s so-called second atmosphere.

An artist’s concept of the early Earth. Image credit: NASA. (Click image for more.)

This is where carbon dioxide enters stage left… er… stage under? Anyway, it was there, right at this early point, along with water vapor, nitrogen, and smaller amounts of other gases. (Note, no oxygen, that is, O2 – significant amounts of that didn’t turn up for another 1.7 billion years, or 2.8 billion years ago.) In fact, carbon dioxide wasn’t just there, it made up most of Earth’s atmosphere, probably not so different from Mars’s atmosphere today.

The point being that carbon dioxide is not a new phenomenon. It is, in fact, the very definition of an old phenomenon. It’s been around, well, pretty much forever. And so has the greenhouse effect. The early Earth was hot. Really hot. Possibly 200 oC or so, because these atmospheric gases trapped the Sun’s heat. Over time, lots and lots of time, the carbon dioxide levels reduced as it became trapped in carbonate rocks, dissolved in the oceans and was utilised by lifeforms for photosynthesis.

Fast-forward a few billion years to the beginning of the twentieth century and atmospheric carbon dioxide levels were about 300 ppm (0.03%), tiny compared to oxygen (about 20%) and nitrogen (about 78%).

Chemists and carbon dioxide

Flemish chemist Jan Baptist van Helmont carried out an experiment which eventually led to the discovery of carbon dioxide gas.

Let’s pause there for a moment and have a little look at some human endeavours. In about 1640 Flemish chemist Jan Baptist van Helmont discovered that if he burned charcoal in a closed vessel, the mass of the resulting ash was much less than that of the original charcoal. He had no way of knowing, then, that he had formed and collected carbon dioxide gas, but he speculated that some of the charcoal had been transmuted into spiritus sylvestris, or “wild spirit”.

In 1754 Scottish chemist Joseph Black noticed that heating calcium carbonate, aka limestone, produced a gas which was heavier than air and which could “not sustain fire or animal life”. He called it “fixed air”, and he’s often credited with carbon dioxide’s discovery, although arguably van Helmont got there first. Black was also the first person to come up with the “limewater test“, where carbon dioxide is bubbled through a solution of calcium hydroxide. He used the test to demonstrate that carbon dioxide was produced by respiration, an experiment still carried out in schools more than 250 years later to show that the air we breathe out contains more carbon dioxide than the air we breathe in.

In 1772 that most famous of English chemists, Joseph Priestley, experimented with dripping sulfuric acid (or vitriolic acid, as he knew it) on chalk to produce a gas which could be dissolved in water. Priestley is often credited with the invention of soda water as a result (more on this in a bit), although physician Dr William Brownrigg probably discovered carbonated water earlier – but he never published his work.

In the late 1700s carbon dioxide became more widely known as “carbonic acid gas”, as seen in this article dated 1853. In 1823 Humphry Davy and Michael Faraday manged to produce liquified carbon dioxide at high pressures. Adrien-Jean-Pierre Thilorier was the first to describe solid carbon dioxide, in 1835. The name carbon dioxide was first used around 1869, when the term “dioxide” came into use.

A diagram from “Impregnating Water with Fixed Air”, printed for J. Johnson, No. 72, in St. Pauls Church-Yard, 1772.

Back to Priestley for a moment. In the late 1800s, a glass of volcanic spring water was a common treatment for digestive problems and general ailments. But what if you didn’t happen to live near a volcanic spring? Joseph Black, you’ll remember, had established that CO2 was produced by living organisms, so it occurred to Priestly that perhaps he could hang a vessel of water over a fermentation vat at a brewery and collect the gas that way.

But it wasn’t very efficient. As Priestly himself said, “the surface of the fixed air is exposed to the common air, and is considerably mixed with it, [and] water will not imbibe so much of it by the process above described.”

It was then that he tried his experiment with vitriolic acid, which allowed for much greater control over the carbonation process. Priestly proposed that the resulting “water impregnated with fixed air” might have a number of medical applications. In particular, perhaps because the water had an acidic taste in a similar way that lemon-infused water does, he thought it might be an effective treatment for scurvy. Legend has it that he gave the method to Captain Cook for his second voyage to the Pacific for this reason. It wouldn’t have helped of course, but it does mean that Cook and his crew were some of the first people to produce carbonated water for the express purpose of drinking a fizzy drink.

Refreshing fizz

You will have noticed that, despite all his work, there is no fizzy drink brand named Priestly (at least, not that I know of).

Joseph Priestley is credited with developing the first method for making carbonated water.

But there is one called Schweppes. That’s because a German watchmaker named Johann Jacob Schweppe spotted Priestley’s paper and worked out a simpler, more efficient process, using sodium bicarbonate and tartaric acid. He went on to found the Schweppes Company in Geneva in 1783.

Today, carbonated drinks are made a little differently. You may have heard about carbon dioxide shortages this summer in the U.K. These arose because these days carbon dioxide is actually collected as a by-product of other processes. In fact, after several bits of quite simple chemistry that add up to a really elegant sequence.

From fertiliser to fizzy drinks

It all begins, or more accurately ends, with ammonia fertiliser. As any GCSE science student who’s been even half paying attention can tell you, ammonia is made by reacting hydrogen with nitrogen during the Haber process. Nitrogen is easy to get hold of – as I’ve already said it makes up nearly 80% of our atmosphere – but hydrogen has to be made from hydrocarbons. Usually natural gas, or methane.

This involves another well-known process, called steam reforming, in which steam is reacted with methane at high temperatures in the presence of a nickel catalyst. This produces carbon monoxide, a highly toxic gas. But no problem! React that carbon monoxide with more water in the presence of a slightly different catalyst and you get even more hydrogen. And some carbon dioxide.

Fear not, nothing is wasted here! The CO2 is captured and liquified for all sorts of food-related and industrial uses, not least of which is fizzy drinks. This works well for all concerned because steam reforming produces large amounts of pure carbon dioxide. If you’re going to add it to food and drinks after all, you wouldn’t want a product contaminated with other gases.

Carbon dioxide is a by-product of fertiliser manufacture.

We ended up with a problem this summer in the U.K. because ammonia production plants operate on a schedule which is linked to the planting season. Farmers don’t usually apply fertiliser in the summer – when they’re either harvesting or about to harvest crops – so many ammonia plants shut down for maintenance in April, May, and June. This naturally leads to reduction in the amount of available carbon dioxide, but it’s not normally a problem because the downtime is relatively short and enough is produced the rest of year to keep manufacturers supplied.

This year, though, natural-gas prices were higher, while the price of ammonia stayed roughly the same. This meant that ammonia plants were in no great hurry to reopen, and that meant many didn’t start supplying carbon dioxide in July, just when a huge heatwave hit the UK, coinciding with the World Cup football (which tends to generate a big demand for fizzy pop, for some reason).

Which brings us back to our atmosphere…

Carbon dioxide calamity?

Isn’t there, you may be thinking, too much carbon dioxide in our atmosphere? In fact, that heatwave you just mentioned, wasn’t that a global warming thing?  Can’t we just… extract carbon dioxide from our air and solve everyone’s problems? Well, yes and no. Remember earlier when I said that at the beginning of the twentieth century and atmospheric carbon dioxide levels were about 300 ppm (0.03%)?

Over the last hundred years atmospheric carbon dioxide levels have increased from 0.03% to 0.04%

Today, a little over 100 years later, levels are about 0.04%. This is a significant increase in a relatively short period of time, but it’s still only a tiny fraction of our atmosphere (an important tiny fraction nonetheless – we’ll get to that in a minute).

It is possible to distill gases from our air by cooling air down until it liquefies and then separating the different components by their boiling points. For example nitrogen, N2, boils at a chilly -196 oC whereas oxygen, O2, boils at a mere -183 oC.

But there’s a problem: CO2 doesn’t have a liquid state at standard pressures. It forms a solid, which sublimes directly into a gas. For this reason carbon dioxide is usually removed from cryogenic distillation mixtures, because it would freeze solid and plug up the equipment. There are other ways to extract carbon dioxide from air but although they have important applications (keep reading) they’re not practical ways to produce large volumes of the gas for the food and drink industries.

Back to the environment for a moment: why is that teeny 0.04% causing us such headaches? How can a mere 400 CO2 molecules bouncing around with a million other molecules cause such huge problems?

For that, I need to take a little diversion to talk about infrared radiation, or IR.

Infrared radiation was first discovered by the astronomer William Herschel in 1800. He was trying to observe sun spots when he noticed that his red filter seemed to get particularly hot. In what I’ve always thought was a rather amazing intuitive leap, he then passed sunlight through a prism to split it, held a thermometer just beyond the red light that he could see with his eyes, and discovered that the thermometer showed a higher temperature than when placed in the visible spectrum.

He concluded that there must be an invisible form of light beyond the visible spectrum, and indeed there is: infrared light. It turns out that slightly more than half of the total energy from the Sun arrives on Earth in the form of infrared radiation.

What has this got to do with carbon dioxide? It turns out that carbon dioxide, or rather the double bonds O=C=O, absorb a lot of infrared radiation. By contrast, oxygen and nitrogen, which make up well over 90% of Earth’s atmosphere, don’t absorb infrared.

CO2 molecules also re-emit IR but, having bounced around a bit, not necessarily in the same direction and – and this is the reason that tiny amounts of carbon dioxide cause not so tiny problems – they transfer energy to other molecules in the atmosphere in the process. Think of each CO2 molecule as a drunkard stumbling through a pub, knocking over people’s pints and causing a huge bar brawl. A single disruptive individual can, indirectly, cause a lot of others to find themselves bruised and bleeding and wondering what the hell just happened.

Like carbon dioxide, water vapour also absorbs infrared, but it has a relatively short lifetime in our atmosphere.

Water vapor becomes important here too, because while O2 and N2 don’t absorb infrared, water vapour does. Water vapour has a relatively short lifetime in our atmosphere (about ten days compared to a decade for carbon dioxide) so its overall warming effect is less. Except that once carbon dioxide is thrown into the mix it transfers extra heat to the water, keeping it vapour (rather than, say, precipitating as rain) for longer and pushing up the temperature of the system even more.

Basically, carbon dioxide molecules trap heat near the planet’s surface. This is why carbon dioxide is described as a greenhouse gas and increasing levels are causing global warming. There are people who are still arguing this isn’t the case, but truly, they’ve got the wrong end of the (hockey) stick.

It’s not even a new concept. Over 100 years ago, in 1912, a short piece was published in the Rodney and Otamatea Times which said: “The furnaces of the world are now burning about 2,000,000,000 tons of coal a year. When this is burned, uniting with oxygen, it adds about  7,000,000,000 tons of carbon dioxide to the atmosphere yearly. This tends to make the air a more effective blanket for the earth and to raise its temperature.”

This summer has seen record high temperatures and some scientists have been warning of a “Hothouse Earth” scenario.

This 1912 piece suggested we might start to see effects in “centuries”. In fact, we’re seeing the results now. As I mentioned earlier, this summer has seen record high temperatures and some scientists have been warning of “Hothouse Earth” scenario, where rising temperatures cause serious disruptions to ecosystems, society, and economies. The authors stressed it’s not inevitable, but preventing it will require a collective effort. They even published a companion document which included several possible solutions which, oddly enough, garnered rather fewer column inches than the “we’re all going to die” angle.

Don’t despair, DO something…

But I’m going to mention it, because it brings us back to CO2. There’s too much of it in our atmosphere. How can we deal with that? It’s simple really: first, stop adding more, i.e. stop burning fossil fuels. We have other technologies for producing energy. The reason we’re still stuck on fossil fuels at this stage is politics and money, and even the most obese of the fat cats are starting to realise that money isn’t much use if you don’t have a habitable planet. Well, most of them. (There’s probably no hope for some people, but we can at least hope that their damage-doing days are limited.)

There are some other, perhaps less obvious, sources of carbon dioxide and other greenhouse gases that might also be reduced, such as livestock, cement for building materials and general waste.

Forests trap carbon dioxide in land carbon sinks. More biodiverse systems generally store more carbon.

And then, we’re back to taking the CO2 out of the atmosphere. How? Halting deforestation would allow more CO2 to be trapped in so-called land carbon sinks. Likewise, good agricultural soil management helps to trap carbon underground. More biodiverse systems generally store more carbon, so if we could try to stop wiping out land and coastal systems, that would be groovy too. Finally, there’s the technological solution: carbon capture and storage, or CSS.

This, in essence, involves removing CO2 from the atmosphere and storing it in geological formations. The same thing the Earth has done for millenia, but more quickly. It can also be linked to bio-energy production in a process known as BECCS. It sounds like the perfect solution, but right now it’s energy intensive and expensive, and there are concerns that BECCS projects could end up competing with agriculture and damaging conservation efforts.

A new answer from an ancient substance?

Forming magnesite, or magnesium carbonate, may be one way to trap carbon dioxide.

Some brand new research might offer yet another solution. It’s another carbon-capture technology which involves magnesium carbonate, or magnesite (MgCO3). Magnesite forms slowly on the Earth’s surface, over hundreds of thousands of years, trapping carbon dioxide in its structure as it does.

It can easily be made quickly at high temperatures, but of course if you have to heat things up, you need energy, which might end up putting as much CO2 back in as you’re managing to take out. Recently a team of researchers at Trent University in Canada have found a way to form magnesite quickly at room temperature using polystyrene microspheres.

This isn’t something which would make much difference if, say, you covered the roof of everyone’s house with the microspheres, but it could be used in fuel-burning power generators (which could be burning renewables or even waste materials) to effectively scrub the carbon dioxide from their emissions. That technology on its own would make a huge difference.

And so here we are. Carbon dioxide is one of the oldest substances there is, as “natural” as they come. From breathing to fizzy drinks to our climate, it’s entwined in every aspect of our everyday existence. It is both friend and foe. Will we work out ways to save ourselves from too much of it in our atmosphere? Personally, I’m optimistic, so long as we support scientists and engineers rather than fight them…


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