More from Genius Lab Gear: Science Word Magnets

Magnets say: the results say we can inhale hot ketones

(Don’t try this at home. Or in the lab.)

The brilliant people at Genius Lab Gear (inventors of The Pocket Chemist) recently sent me a new toy: Science Word Magnets!

They are, as the name suggests, magnetic words, but with the twist that they have science and engineering themes. There are sets for ecology, engineering, microbiology, neuroscience, physics and, of course, chemistry. There’s also a science basics set, an academia set and a PhD balance set.

I’ve been messing about with the science basics set, the starter tile set ($3 extra with any order) and the chemistry set, and they really are loads of fun!

Board shows random magnets

These science word magnets have been specially designed by experts in each field to have technical depth while being fun to use.

Stick them on your fridge, your magnetic whiteboard, or anywhere you might usually persuade a magnet to stick.

And guess what? Yes, there’s a discount code! Use FLASKMAG1 when you check out to save $1 on each set you buy (so the more you buy, the more you save).

magnets read: question, method, experiment, scientific notebook, equations, formulas, results, publish, tequila

The magnets fit with other popular word magnet sets.

Follow this link and the code will be automatically applied.

By doing so, you’ll also be supporting this site, and helping to fund more cool chemistry articles — thank you!

Shipping is FREE for the USA and Canada (no tracking) and $5.90 for the UK, Europe, Japan, Korea and Australia. Shipping for elsewhere in the world is calculated at checkout. Add 4 sets to get $5 OFF and free expedited shipping in the USA!


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.

Lovely lollipops: the chemistry of sugary things

20th July is National Lollipop Day!

Today, July 20th, is apparently national lollipop day in the United States, and general news is… *waves hands* so it seems like a good excuse to write something with lots of pictures of brightly coloured sweets, right? Plus, sugar!

The idea of putting something sugary on a stick to hold and eat is an ancient one. The very earliest humans probably used sticks to collect honey from beehives. Later, the Chinese, Egyptians and people from the Middle East dipped fruits and nuts in honey and used sticks to make them easier to eat.

In the 17th century, boiled sugar sweets were made in England and, again, sticks inserted to make eating easier. This may be where the name “lollipop” originates, since “lolly” is a dialect word for tongue. Later, in the American Civil War era (early 1860s), some sources say hard candy was put on the tips of pencils for children. In 1931 an American named George Smith started making hard candies on sticks, and trademarked the name lollipop — but he reportedly took the name from a racehorse named “Lolly Pop”.

Table sugar is sucrose

Enough history, let’s get to the chemistry! Lollipops are made of sugar, with added colours and flavours. I’ve talked about sugar before, and it’s always worth remembering that we tend to use the word rather loosely in everyday speech.

There’s more than one type of sugar: in particular, the three that are probably most familiar are glucose, fructose and sucrose. Glucose is a simple sugar, and the one you might remember from photosynthesis and respiration equations. It’s essential for life, and you quickly run into serious trouble if your blood glucose levels drop too low (just ask a diabetic).

Like glucose, fructose is a monosaccharide (the simplest form of sugar), and is often called “fruit sugar” because, guess what, it’s common in fruits. Sucrose is what we know as “table sugar” and is a disaccharide, made up of a unit of glucose joined to a unit of fructose. In the body, sucrose is broken up into glucose and fructose.

Rock candy is made from sucrose but, unlike in most lollipops and hard candy, the sugar is allowed to form large crystals

The primary ingredient in lollipops is usually sucrose, which can be persuaded (more in a minute) to set nicely to produce a hard, shiny surface. However, commercial lollipops often also include corn syrup, or glucose syrup, which contains oligosaccharides: larger sugar molecules made from a number of simple sugar molecules joined together. Typically, as the name “glucose syrup” might suggest, these molecules contain units of glucose.

It’s worth mentioning here that corn syrup/glucose syrup isn’t the same as “high fructose corn syrup” or HFCS, in which the glucose molecules have been converted into fructose. This product is cheap, sweet and commercially easy to use, but it’s also controversial. Excessive consumption has been linked to obesity and non-alcoholic fatty liver disease, although the actual evidence is weak: a systematic review in 2014 concluded that there was little evidence it was worse than other forms of sugar. It’s really a problem of quantity: it’s easy and cheap for food manufacturers to throw HFCS into foods and drinks, and of course it tastes delicious, so as a consequence consumers end up eating too much of the stuff. In short: more water and fruit, less cake and fizzy drinks.

But having done the obligatory “eat healthily” thing, one lollipop isn’t going to hurt, is it? So back to that…

Fudge, perhaps surprisingly, contains the crystalline form of sugar

When it cools, sugar forms two different types of solid: crystalline and glassy amorphous (sometimes described as ‘amorphous solid’). Now, you might imagine that sugar as a crystalline solid is found in hard sweets/candies, but, no — it mostly turns up in soft things like fudge and fondant, which contain lots of very tiny crystals, giving an ever-so slightly granular texture. (An exception is rock candy, where the sugar is encouraged to form large crystals.)

The glassy amorphous form of sugar, on the other hand, can be literally like glass: hard, brittle, and transparent. In fact, “sugar glass” has in the past been used to make windows, bottles and so on for special effects in film and television, because it’s much less likely to cause injury than “real” glass. However, it’s very fragile and hygroscopic (meaning it absorbs water, causing it to soften over time) so these days it’s largely been replaced by synthetic resins.

Honey can be used as an inhibitor, to prevent crystallisation

The glassy amorphous form of sugar is achieved by starting with a 50% sugar solution which also contains an inhibitor, to prevent crystals forming spontaneously. Common inhibitors are the corn syrup I mentioned earlier, or cream of tartar (potassium bitartrate), honey or butter.

Exactly which you use depends on the recipe, but they all do essentially the same thing, namely, get in the way of the glucose molecules and prevent them ordering themselves into a regular (crystalline) structure. The mixture is heated to a high temperature (about 155 oC) until almost all the water evaporates — the final candy will only have about 1-2% water — and then cooled until glass transition occurs.

At the glass transition point, the sugar mixture becomes solid.

This is the clever bit, and only happens if crystallisation is inhibited (else crystals form instead). Glass transition happens around 100-150 oC below the melting point of the pure substance. For example, the melting point of pure sucrose is 186 oC, but it undergoes glass transition at around 60 oC.

Glass transition is a reversible change, which we might (if I didn’t generally dislike the concept) call a physical change. It’s a change of phase, where the sugar mixture changes from liquid to solid, but it’s different from crystallisation, because instead of the molecules becoming more ordered, they simply ‘freeze’ in their random, liquid positions. (It is, for the record, annoyingly difficult to show this in diagram form.)

Amorphous solid structures are sometimes called “supercooled liquids”. This isn’t wrong, but personally I think it’s unhelpful (and can lead to nonsense about glass flowing very slowly over time). Once cooled and set, glass, whether window glass or sugar glass, is absolutely not a liquid; it’s a solid.

Of course, to make lollipops, all sorts of colours and flavours are added to the mixture as well, and sometimes more than one mixture is used to create intricate, layered effects. There are even medicinal lollipops which contain, for example, the powerful painkiller fentanyl — the idea being that the patient can administer the dose gradually as needed.

Which brings me to the end. Happy National Lollipop Day! My favourites are Chupa Chups — if you’ve enjoyed this, how about popping over to Ko-fi so I can stock up? And if you’ve been eating sweets, do remember to clean your teeth!


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.

Chemical connections: dexamethasone, hydroxychloroquine and rheumatoid arthritis

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 ReichsteinEdward 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!)

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

Easy Indicators

Indicator rainbow, reproduced with kind permission of Isobel Everest, @CrocodileChemi1

Recently on Twitter CrocodileChemist (aka Isobel Everest), a senior school science technician (shout out to science technicians, you’re all amazing) shared a fabulous video and photo of a “pH rainbow”.

The effect was achieved by combining various substances with different pH indicators, that is, substances that change colour when mixed with acids or alkalis.

Now, this is completely awesome, but, not something most people could easily reproduce at home, on account of their not having methyl orange or bromothymol blue, or a few other things (that said, if you did want to try, Isobel’s full method, and other indicator art, can be found here).

But fear not, I’ve got this. Well, I’ve got a really, really simple version. Well, actually, I’ve got more of an experiment, but you could make it into more of a rainbow if you wanted. Anyway…

This is what you need:

  • some red cabbage (one leaf is enough)
  • boiling water
  • mug
  • white plate, or laminated piece of white card, or white paper in a punched pocket
  • cling film/clear plastic wrap (if you’re using a plate)
  • mixture of household substances (see below)
  • board marker (optional) or pen
  • plastic pipettes (optional, but do make it easier – easily bought online)

First, make the indicator. There are recipes online, but some of them are over-complicated. All you really need to do is finely chop the red cabbage leaf, put it in a mug, and pour boiling water over it. Leave it to steep and cool down. Don’t accidentally drink it thinking it’s your coffee. Pour off the liquid. Done.

If you use a plate, cover it with cling film

Next, if you’re using a plate, cover it with cling film. There are two reasons for this: firstly, cling film is more hydrophobic (water-repelling) than most well-washed ceramic plates, so you’ll get better droplets. Secondly, if you write on a china plate with a board marker it doesn’t always wash off. Ask me how I know.

Next step: hunt down some household chemicals. I managed to track down oven cleaner, plughole sanitiser, washing up liquid, lemon juice, vinegar, limescale remover and toilet cleaner (note: not bleach – don’t confuse these two substances, one is acid, one is alkali, and they must never be mixed).

Label your plate/laminated card/paper in punched pocket with the names of the household substances.

Place a drop of cabbage indicator by each label. Keep them well spaced so they don’t run into each other. Also, at this stage, keep them fairly small. Leave one alone as a ‘control’. On my plate, it’s in the middle.

Add a drop of each of your household substances and observe the colours!

Red cabbage indicator with various household substances

IMPORTANT SAFETY NOTE: some of these substances are corrosive. The risk is small because you’re only using drops, but if working with children, make sure an adult keeps control of the bottles, and they only have access to a tiny amount. Drip the more caustic substances yourself. Take the opportunity to point out and explain hazard warning labels. Use the same precautions you would use when handling the substance normally, i.e. if you’d usually wear gloves to pick up the bottle, wear gloves. Some of these substances absolutely must not be mixed with each other: keep them all separate.

Here’s a quick summary of what I used:

A useful point to make here is that pH depends on the concentration of hydrogen ions (H+) in the solution. The more hydrogen ions, the more acidic the solution is. In fact, pH is a log scale, which means a change of x10 in hydrogen concentration corresponds to a change of one pH point. In short, the pH of a substance changes with dilution.

Compound Interest’s Cabbage Indicator page (click image for more info)

Which means that if you add enough water to acid, the pH goes up. So, for example, although the pH of pure ethanoic acid is more like 2.4, a dilute vinegar solution is probably closer to 3, or even a bit higher.

Compound Interest, as is usually the case, has a lovely graphic featuring red cabbage indicator. You can see that the colours correspond fairly well, although it does look like my oven cleaner is less alkaline (closer to green) than the plughole sanitiser (closer to yellow).

As the Compound Interest graphic mentions, the colour changes are due to anthocyanin pigments. These are red/blue/purple pigments that occur naturally in plants, and give them a few advantages, one of which is to act as a visual ripeness indicator. For example, the riper a blackberry is, the darker it becomes. That makes it stand out against green foliage, so it’s easier for birds and animals to find it, eat it and go on to spread the seeds. Note that “unripe” colours, yellow-green, are at the alkaline end, which corresponds to bitter flavours. “Ripe” colours, purple-red, are neutral to acidic, corresponding with much more appealing sweet and tart flavours. Isn’t nature clever?

You can make a whole mug full of indicator from a single cabbage leaf (don’t drink it by mistake).

Which brings me to my final point – what if you can’t get red cabbage? Supermarkets are bit… tricky at the moment, after all. Well, try with some other things! Any dark-coloured plant/fruit should work. Blueberries are good (and easy to find frozen). The skins of black grapes or the very dark red bit of a rhubarb stalk are worth a try. Blackberries grow wild in lots of places later in the year. Tomatoes, strawberries and other red fruits will also give colour changes (I’ve talked about strawberries before), although they’re less dramatic.

For those (rightly) concerned about wasting food – you don’t need a lot. I made a whole mug full of cabbage indicator from a single cabbage leaf, and it was the manky brown-around-the-edges one on the outside that was probably destined for compost anyway.

So, off you go, have fun! Stay indoors, learn about indicators, and stay safe.

EDIT: after I posted this, a few people tried some more experiments with fruits, vegetables and plants! Beaulieu Biology posted the amazing grid below, which includes everything from turmeric to radishes:

Image reproduced with kind permission of Beaulieu Biology (click for larger version)

And Compound Interest took some beautiful photos of indicator solutions extracted from a tulip flower, while CrocodileChemist did something similar and used the solutions to make a gorgeous picture of a tree. Check them out!


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

Cleaning chemistry – the awesome power of soap

Well, times are interesting at the moment, aren’t they? I’m not going to talk (much) about The Virus (there’s gonna be a movie, mark my words), because everyone else is, and I’m not an epidemiologist, virologist or an immunologist or, in fact, in any way remotely qualified. I am personally of the opinion that it’s not even especially helpful to talk about possibly-relevant drugs at the moment, given that we don’t know enough about possible negative interactions, and we don’t have reliable data about the older medicines being touted.

In short, I think it’s best I shut up and leave the medical side to the experts. But! I DO know about something relevant. What’s that, I hear you ask? Well, it’s… soap! But wait, before you start yawning, soap is amazing. It is fascinating. It both literally and figuratively links loads of bits of cool chemistry with loads of other bits of cool chemistry. Stay with me, and I’ll explain.

First up, some history (also not a historian, but that crowd is cool, they’ll forgive me) soap is old. Really, really, old. Archaeological evidence suggests ancient Babylonians were making soap around 4800 years ago – probably not for personal hygiene, but rather, mainly, to clean cooking pots. It was originally made from fats boiled with ashes, and the theory generally goes that the discovery was a happy accident: ashes left from cooking fires made it much easier to clean pots and, some experimenting later, we arrived at something we might cautiously recognise today as soap.

Soap was first used to clean pots.

The reason this works is that ashes are alkaline. In fact, the very word “alkali” is derived from the Arabic al qalīy, meaning calcined ashes. This is because plants, and especially wood, aren’t just made up of carbon and hydrogen. Potassium and calcium play important roles in tree and plant metabolism, and as a result both are found in moderately significant quantities in wood. When that wood is burnt at high temperatures, alkaline compounds of potassium and calcium form. If the temperature gets high enough, calcium oxide (lime) forms, which is even more alkaline.

You may, in fact, have heard the term potash. This usually refers to salts that contain potassium in a water-soluble form. Potash was first made by taking plant ashes and soaking them in water in a pot, hence, “pot ash”. And, guess where we get the word potassium from? Yep. The pure element, being very reactive, wasn’t discovered until 1870, thousands of years after people first discovered how useful its compounds could be. And, AND, why does the element potassium have the symbol K? It comes from kali, the root of the word alkali.

See what I mean about connections?

butyl ethanoate butyl ethanoate

Why is the fact that the ashes are alkaline relevant? Well, to answer that we need to think about fats. Chemically, fats are esters. Esters are chains of hydrogen and carbon that have, somewhere within them, a cheeky pair of oxygen atoms. Like this (oxygen atoms are shown red):

Now, this is a picture of butyl ethanoate (aka butyl acetate – smells of apples, by the way) and is a short-ish example of an ester. Fats generally contain much longer chains, and there are three of those chains, and the oxygen bit is stuck to a glycerol backbone.

Thus, the thick, oily, greasy stuff that you think of as fat is a triglyceride: an ester made up of three fatty acid molecules and glycerol (aka glycerine, yup, same stuff in baking). But it’s the ester bit we want to focus on for now, because esters react with alkalis (and acids, for that matter) in a process called hydrolysis.

Fats are esters. Three fatty acid chains are attached to a glycerol “backbone”.

The clue here is in the name – “hydro” suggesting water – because what happens is that the ester splits where those (red) oxygens are. On one side of that split, the COO group of atoms gains a metal ion (or a hydrogen, if the reaction was carried out under acidic conditions), while the other chunk of the molecule ends up with an OH on the end. We now have a carboxylate salt (or a carboxylic acid) and an alcohol. Effectively, we’ve split the molecule into two pieces and tidied up the ends with atoms from water.

Still with me? This is where it gets clever. Having mixed our fat with alkali and split our fat molecules up, we have two things: fatty acid salts (hydrocarbon chains with, e.g. COONa+ on the end) and glycerol. Glycerol is extremely useful stuff (and, funnily enough, antiviral) but we’ll put that aside for the moment, because it’s the other part that’s really interesting.

What we’ve done here is produce a molecule that has a polar end (the charged bit, e.g. COONa+) and a non-polar end (the long chain of Cs and Hs). Here’s the thing: polar substances tend to only mix with other polar substances, while non-polar substances only mix with other non-polar substances.

You may be thinking this is getting technical, but honestly, it’s not. I guarantee you’ve experienced this: think, for example, what happens if you make a salad dressing with oil and vinegar (which is mostly water). The non-polar oil floats on top of the polar water and the two won’t stay mixed. Even if you give them a really good shake, they separate out after a few minutes.

The dark blue oily layer in this makeup remover doesn’t mix with the watery colourless layer.

There are even toiletries based around this principle. This is an eye and lip makeup remover designed to remove water-resistant mascara and long-stay lipstick. It has an oily layer and a water-based layer. To use it, you give the container a good shake and use it immediately. The oil in the mixture removes any oil-based makeup, while the water part removes anything water-based. If you leave the bottle for a minute or two, it settles back into two layers.

But when we broke up our fat molecules, we formed a molecule which can combine with both types of substance. One end will mix with oily substances, and the other end mixes with water. Imagine it as a sort of bridge, joining two things that otherwise would never be connected (see, literal connections!)

There are a few different names for this type of molecule. When we’re talking about food, we usually use “emulsifier” – a term you’ll have seen on food ingredients lists. The best-known example is probably lecithin, which is found in egg yolks. Lecithin is the reason mayonnaise is the way it is – it allows oil and water to combine to give a nice, creamy product that stays mixed, even if it’s left on a shelf for months.

When we’re talking about soaps and detergents, we call these joiny-up molecules “surfactants“. You’re less likely to have seen that exact term on cosmetic ingredients lists, but you will (if you’ve looked) almost certainly have seen one of the most common examples, which is sodium laureth sulfate (or sodium lauryl sulfate), because it turns up everywhere: in liquid soap, bubble bath, shampoo and even toothpaste.

I won’t get into the chemical makeup of sodium laureth sulfate, as it’s a bit different. I’m going to stick to good old soap bars. A common surfactant molecule that you’ll find in those is sodium stearate, which is just like the examples I was talking about earlier: a long hydrocarbon chain with COONa+ stuck on the end. The hydrocarbon end, or “tail”, is hydrophobic (“water-hating”), and only mixes with oily substances. The COONa+ end, or “head”, is hydrophilic (“water loving”) and only mixes with watery substances.

Bars of soap contain sodium stearate.

This is perfect because dirt is usually oily, or is trapped in oil. Soap allows that oil to mix with the water you’re using to wash, so that both the oil, and anything else it might be harbouring, can be washed away.

Which brings us back to the wretched virus. Sars-CoV-2 has a lipid bilayer, that is, a membrane made of two layers of lipid (fatty) molecules. Virus particles stick to our skin and, because of that membrane, water alone does a really bad job of removing them. However, the water-hating tail ends of surfacant molecules are attracted to the virus’s outer fatty surface, while the water-loving head ends are attracted to the water that’s, say, falling out of your tap. Basically, soap causes the virus’s membrane to dissolve, and it falls apart and is destroyed. Victory is ours – hurrah!

Hand sanitisers also destroy viruses. Check out this excellent Compound Interest graphic (click the image for more).

Who knew a nearly-5000 year-old weapon would be effective against such a modern scourge? (Well, yes, virologists, obviously.) The more modern alcohol hand gels do much the same thing, but not quite as effectively – if you have access to soap and water, use them!

Of course, all this only works if you wash your hands thoroughly. I highly recommend watching this video, which uses black ink to demonstrate what needs to happen with the soap. I thought I was washing my hands properly until I watched it, and now I’m actually washing my hands properly.

You may be thinking at this point (if you’ve made it this far), “hang on, if the ancient Babylonians were making soap nearly 5000 years ago, it must be quite easy to make… ooh, could I make soap?!” And yes, yes it is and yes you can. Believe me, if the apocolypse comes I shall be doing just that. People rarely think about soap in disaster movies, which is a problem, because without a bit of basic hygine it won’t be long before the hero is either puking his guts up or dying from a minor wound infection.

Here’s the thing though, it’s potentially dangerous to make soap, because most of the recipes you’ll find (I won’t link to any, but a quick YouTube search will turn up several – try looking for “saponification“) involve lye. Lye is actually a broad term that covers a couple of different chemicals, but most of the time when people say lye these days, they mean pure sodium hydroxide.

Pure sodium hydroxide is usually supplied as pellets.

Pure sodium hydroxide comes in the form of pellets. It’s dangerous for two reasons. Firstly, precisely because it’s so good at breaking down fats and proteins, i.e. the stuff that humans are made of, it’s really, really corrosive and will give you an extremely nasty burn. Remember that scene in the movie Fight Club? Yes, that scene? Well, that. (Follow that link with extreme caution.)

And secondly, when sodium hydroxide pellets are mixed with water, the solution gets really, really hot.

It doesn’t take a lot of imagination to realise that a really hot, highly corrosive, solution is potentially a huge disaster waiting to happen. So, and I cannot stress this enough, DO NOT attempt to make your own soap unless you have done a lot of research AND you have ALL the appropriate safety equipment, especially good eye protection.

And there we are. Soap is ancient and awesome, and full of interesting chemistry. Make sure you appreciate it every time you wash your hands, which ought to be frequently!

Stay safe, everyone. Take care, and look after yourselves.


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 poem. Enjoy!

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, please consider buying me a coffee through Ko-fi using the button below.
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Genius Lab Gear: The Pocket Chemist

The lovely people at Genius Lab Gear were kind enough to send me one of these to try the other day: The Pocket Chemist!

The Pocket Chemist is a handy double-sided stencil and chemistry reference.

It’s a double-sided stencil which is also printed with lots of really useful chemistry reference information.

It’s made of enamel-coated stainless steel, which not only gives it a really solid, quality feel, but also means you can spill acetone on it without fear.

The edges are super-straight, so you can use it as a (85 mm) ruler. It’s marked in inches and centimetres, includes a small protractor for measuring angles, and there are stencils for various cyclic compounds—including a hexagon so your benzene rings will always be immaculate.

On the back, there’s a full (if small) periodic table that, yes, has the correct symbols for the four elements that were last to get their names (if your eyes are struggling, click on the photo to see a bigger vision).

There’s a full periodic table on the back (click on the image for a larger version).

There’s plenty of other useful information, too: formulas for pH calculations, Gibbs free energy change and others, a number of useful constants (including Avogadro’s number and the molar gas constant in three different unit forms) and other handy bits and pieces such as prefixes for large and small numbers.

Another clever feature is a phone stand slot: put a sturdy credit card-sized card in the straight line at the top, and you can use it to rest your phone at an angle. It’s not strong enough for heavy-handed screen-jabbing, but it works well enough if you just want to watch a video.

Use the stencils to ensure your hexagons are always perfect!

I have to say, I genuinely love the Pocket Chemist. What a great idea. It’s well-made and the perfect size to fit into your wallet, pocket or pencil case. It’s the perfect piece of kit to take to lessons or lectures (no sneaking it into exams, though!).

Now for the good bit: I’ve got a discount code for you! Order from Genius Lab Gear and enter the code FLASK15 at check out, and you’ll get 15% off your order (and I get a small commission which helps pay for this site—win, win!). Shipping is FREE.

Quick note for my non-American readers: with a few minor exceptions, shipping is free worldwide (it’s a thin item that fits in a regular envelope) and delivery is pretty quick.

… AND if you’d like some Science Word Magnets from the same people, check out this page for a discount code for those!


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Blue skies and copper demons: a story of mysterious purple crystals

Mystery purple crystals (posted with permission of Caroline Hedge, @CM_Hedge)

Today, a little story about some mysterious, purple crystals. On Tuesday, Twitter user Caroline Hedge posted this photo with the question: “What the %#&$ is lab putting down the drain to cause this?”

The post spawned lots of responses, some more serious than others. One of the sensible ones came from Roland Roesler, who thought that the pipe had corroded from the outside, suggesting that a leaky connection at the top right had allowed sewage to drip down the right-hand side of the copper pipe and drip from the bottom, which explained why the left-hand half of the pipe appeared unscathed.

I agreed. The pipe is clearly made of copper, and blue colours are characteristic of hydrated copper salts. Inside the pipe, the flow of water would wash any solution anyway before corrosion could occur, but on the outside, drips could sit on the surface for long periods of time. There’d be plenty of time for even a slow reaction to occur, and then for water to slowly evaporate, allowing the growth of spectacular crystals.

Hydrated copper(II) sulfate crystals are bright blue. (Image from Wikimedia Commons)

But what exactly where they? There were several theories, but for me the interesting thing was the colour. Hydrated copper(II) sulfate crystals are bright blue. The colour arises due to an effect called d orbital splitting, which is a tad complicated but, in short, means that complex absorbs light from the red end of the visible light spectrum, allowing all the other colours of light to pass through. As a result, our eyes “see” blue.

But these crystals, assuming it’s not a photographic effect, had a purplish hue. At least, some of them do. So… not copper sulfate, or not entirely copper sulfate (given the situation, a mixture seemed entirely likely). Which begs the question, which copper complex produces a purple colour?

A little bit of Googling and I was pretty sure I’d identified it: copper azurite, Cu₃(CO₃)₂(OH)₂. This fit for two reasons: firstly, it’s a mineral that could (does) readily form in the presence of water and air (which, of course, contains carbon dioxide), and secondly it’s exactly the right colour.

Many will recognise the word “azure” as being associated with the deep, rich blue of a summer sky, and in fact the English name of this mineral comes from the same word-root: the Persian lazhward, a place known for its deposits of another deep-blue stone, lapis lazuli (meaning “stone of azure”).

Blue-purple copper azurite and green malachite (image from Wikipedia)

Azurite is often found with malachite, the better-known green copper mineral that we recognise from copper roofs and statues. Malachite is sometimes simplistically described as copper carbonate, implying CuCO₃, but in truth it’s Cu₂CO₃(OH)₂ pure copper(II) carbonate doesn’t form in nature.

You can see malachite co-existing with azurite in the photo on the right. The azurite will, over time, tend to morph into malachite when the level of carbon dioxide in the air is relatively low, as in ‘normal’ air—which explains why we don’t usually see purple ‘copper’ roofs—but the carbon dioxide levels were probably higher in that cupboard. There was almost certainly acidic sewage reacting with carbonate, combined with a lack of ventilation, so it makes sense that we might see more azurite.

Azurite has an interesting history as a pigment. Historically blue colours were rare and expensive—associated with royalty and divinity—which is one reason why the Virgin Mary was often depicted wearing blue in paintings. Azurite was used to make blue pigments, but (as I mentioned above) it’s unstable, tending to turn greenish over time, or black if heated. Ultramarine blue (made from lapis lazuli) is more stable, particularly when heated, but it was even more expensive. A lot of blue pigments in medieval paintings have been misidentified as coming from lapis lazuli, when in fact they were azurite—a more common mineral in Europe at the time.

There’s a fun piece of etymology here, too. Copper, of course, has been valuable metal since, well, the Bronze Age. The presence of purple azurite and green malachite are surface indicators of copper sulfide ores, useful for smelting. This lead to the name of the element nickel, because an ore of nickel weathers to produce a green mineral that looks a little like malachite. And this, in turn, lead to attempts to smelt it in the belief that it was copper ore. But, since it wasn’t, the attempts to produce copper failed (a much higher smelting temperature is needed to produce nickel).

The mineral nickeline can resemble malachite, and was dubbed kupfernickel in Germany, literally “copper demon”

As a result, the mineral, nickeline, was dubbed kupfernickel in Germany, literally “copper demon”. When the Swedish alchemist Baron Axel Fredrik Cronstedt succeeded, in 1751, in smelting kupfernickel to produce a previously unknown silvery-white, iron-like metal he named it after the nickel part of kupfernickel.

And this is how we go from a corroded pipe to sky-blue colours to medieval paintings to copper demons to nickel. But what happened to the pipe in the original tweet? Well, in an update, Caroline Hedge told us that it had been removed and disposed of, and so we’ll never be completely sure what the pretty crystals were, but they certainly lead to an interestingly twisty-turny chemistry story.


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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…


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.
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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 voilà, 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 £1 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!


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