Neem: nice, nasty or… not sure?

Posted on 20 Aug, 2021 by katlday

A few days ago it was sunny and slightly breezy outside (yes, it’s August, but I live in the UK – this isn’t as common as you might imagine) and I thought, I should make the most of this and do something about my orchids.

Now, anyone that reads this blog regularly will know that my Dad is a horticulturist. I, however, am not. My fascination with bright colours, interesting smells and complicated naming conventions went down the chemistry route. But I am, oddly, quite good with Phalaenopsis, aka, moth orchids. I don’t really know why, or how, but I seem to have come to some sort of agreement with the ones that live on my kitchen windowsill. It goes along the lines of: I’ll water you once a week, and you make flowers a couple of times a year, and we’ll otherwise leave each other alone, okay?

Scale bugs secrete honeydew, which encourages the growth of sooty moulds

Well, this was fine for years, until we somehow acquired an infestation of scale bugs. These tiny but extremely annoying pests feed by sucking sap from leaves of plants, and they excrete a sticky substance called honeydew. Trust me, it’s not as nice as it sounds. Firstly, it really is sticky, and makes a horrible mess not just of the orchid leaves, but also the area around the plants.

Then it turns out that certain types of mould just love this stuff, so you end up with black spots on the leaves. And, not surprisingly, all this weakens the plant.

So, what’s the answer? Well, there are several. But the one I tend to default to is neem oil.

This stuff is a vegetable oil from the seeds of the Azadirachta indica, or neem, tree. It has a musty, nutty sort of smell, and is fairly easy to buy.

It’s indigenous to the Indian subcontinent and has been historically important in traditional medicine. In fact, The Sanskrit name of this evergreen tree is ‘Arishtha’, which means ‘reliever of sickness’.

So it’s a natural vegetable oil and people have been using it as a remedy for thousands of years – must be totally safe, right? Right?

Well… I’ve said it before, but some of the very best horribly toxic things are entirely natural, and neem is yet another example. Ingestion of significant quantities can cause metabolic acidosis (finally, something that really does have the potential to change blood pH! Er… but not… in a good way), kidney failure, seizures, and brain damage in children. Skin contact can cause contact dermatitis. It’s been shown to work as a contraceptive and, more problematically, it’s a known abortifacient (causes miscarriage).

Neem oil is easy to buy, but it needs to be handled with caution

All this said, as always, the dose make the poison.

One case study in the Journal of The Association of Physicians of India reported on a 36-year-old man who swallowed 30–50 ml (about three tablespoons) of neem oil, in the hope of treating the corns on his feet. As far as I can tell, it didn’t help his corns. It did cause vomiting, drowsiness, a dangerous drop in blood pH and seizures. There’s no specific antidote for neem poisoning, but the hospital managed his symptoms. Luckily, despite the hammering his kidneys undoubtedly took, he didn’t need dialysis, and was discharged from hospital after just over a week.

Now, okay, you’re unlikely to accidentally swallow three tablespoons of any oil, especially not neem which does have quite a strong, not entirely pleasant, smell and (reportedly – I haven’t tried for obvious reasons) a bitter taste. But nevertheless, it’s wise to be cautious, particularly around children who have a smaller body mass and therefore are much more likely to suffer serious effects – up to and including death. In one reported case, a mother gave a 3-month-old child a teaspoon of neem oil in the hope of curing his indigestion – fortunately he survived, but not without some seriously scary symptoms.

Nimbin, a chemical found in neem oil, is reported to have all sorts of beneficial effects [image source]

Okay, so those are the dangers. Let’s talk chemistry. The Pakistani organic chemist Salimuzzaman Siddiqui is thought to be the first scientist to formally investigate the various compounds in neem oil. In 1942 he extracted three compounds, and identified nimbidin as the main antibacterial substance in neem. He was awarded an OBE in 1946 for his discoveries.

I will confess, at this point, to running into a little bit of confusion with the nomenclature. Nimbidin is described, in some places at least, as a mixture of compounds (collectively, tetranortriterpenes) rather than a single molecule. But either way, it has been shown to have anti-inflammatory properties – at least in rats.

Another of the probably-mostly-good substances in neem is nimbin: a triterpenoid which is reported to have a whole range of positive properties, including acting as an anti-inflammatory, an antipyretic, a fungicide, an antiseptic and even as an antihistamine. Interestingly, I went looking for safety data on nimbin, and I couldn’t find much. That could mean it’s safe, or it could mean it just hasn’t been extensively tested.

Azadirachtin, another chemical found in neem, is a known pesticide [image source]

The substance that seems to do most of the pesticide heavy lifting is azadirachtin. This is a limonoid (compounds that are probably best known for their presence in citrus fruits). It’s what’s called an antifeedant – a substance produced by plants to deter predators from munching on them. Well, mostly. Humans have a strange habit of developing a taste for plants that produce such substances. Take, for example, odoriferous garlic, clears-out-your-sinuses horseradish, and of course the daddy of them all: nicotine.

Azadirachtin is known to affect lots of species of insects, both by acting as an antifeedant and as a growth disruptor. Handily, it’s also biodegradable – and breaks down in a few days when exposed to light and water.

That makes it appealing as a potential pesticide, and it’s also generally described as having low toxicity in mammals – its reported LD₅₀ tends to fall into the grams per kilogram range, which makes it “moderately to slightly toxic“. Wikipedia quotes a value (without a source, as I write this) of >3,540 mg/kg in rats.

But… I did find another page quoting 13 mg/kg in mice. That’s quite dramatically different, and would make it extremely/highly toxic. Unfortunately I couldn’t get my hands on the original source, so I haven’t been able to verify it’s not a transposition error.

Let’s assume it isn’t. It would be odd to have such a big difference between mice and rats. Things that poison mice tend to poison rats, too. There might be some confusion over pure azadirachtin vs. “neem extract” – it could be the case that the mixture of chemicals working together in neem create some sort of synergistic (toxic) effect – greater than the sum of all the individual substances. It could be an experimental error, including a contaminated neem sample, or something to do with the way the animals were exposed to the extract.

Neem soap is widely available online, but that may not be a good thing…

It’s difficult to say. Well, it’s difficult for me to say, because I don’t have access to all the primary sources. (Any toxicologists out there, please do feel free to weigh into the comments section!) But either way, as I’ve already mentioned, several case studies have fingered azadirachtin as one of the substances likely to be causing the well-reported nasty side effects.

If you’re asking this chemist? I say be careful with the stuff. If you decide to use it on your plants, keep it out of reach of children, and wear some good-quality disposable gloves while you’re handling it (I put some on after I took that photo back there). If you’re pregnant, or trying to become pregnant, the safest option is to not use it at all.

Which brings me to neem soap.

Yup. It’s sold as a “natural” treatment for skin conditions like acne. I won’t link to a specific brand, but it’s easy to find multiple retailers with a simple Google search. I looked at one selling soap bars for £6.99 a pop, containing 10% (certified organic, because of course) neem oil. Did I mention back there that neem is known to cause contact dermatitis? I’m fairly sure I did. None of the products I saw had obvious safety warnings, and I certainly found nothing about safety (or otherwise) for pregnant women.

Plus – worryingly, not least because children are more likely to get things in their mouth – you can also buy kids and babies versions, again purporting to contain 10% neem oil.

I even found neem toothpaste. Which… given people often swallow toothpaste… yikes.

My moth orchids are looking much healthier now I’ve got rid of all the scale bugs!

Now again, and for the umpteenth time, the dose makes the poison. The case studies I’ve mentioned involved, at a minimum, swallowing a teaspoon of pure neem oil, and you’re not getting that sort of quantity from smears of toothpaste. But, at the same time, when it comes to pregnancy and babies, it’s generally sensible to apply a precautionary principle, especially for things like soap and toothpaste for which alternatives with well-established safety profiles exist.

Bottom line? Would I use these products? I would not.

But I do use neem to treat the scale bugs on my orchids, and they’re doing much better than they were. Fingers crossed for more flowers!

Do you want something non-sciency to distract you from, well, everything? Why not take a look at my fiction blog: the fiction phial? You can also find me doing various flavours of editor-type-stuff at the horror podcast, PseudoPod.org – so head over there, too!

Like the Chronicle Flask’s Facebook page for regular updates, or follow @chronicleflask on Twitter. Content is © Kat Day 2021. You may share or link to anything here, but you must reference this site if you do. You can support my writing my buying a super-handy Pocket Chemist from Genius Lab Gear using the code FLASK15 at checkout (you’ll get a discount, too!) or by buying me a coffee – the button is right here…

Faking Lateral Flow Tests: the problem with pH

Posted on 6 Jul, 2021 by katlday

Fruit juices can be used to generate a fake positive on COVID-19 LFTs

On Thursday last week, I got a message from Prof Mark Lorch — my sometime collaborator on supercharacter-based ramblings.

“Have you seen the reports of kids fooling the Covid lateral flow tests and getting false +ve results by adding orange juice to the devices?” he wondered.

At this point, I had not – but I quickly got up to speed. Mark had previously made an excellent video explaining how lateral flow test (LFT) devices work, so it was just a case of working out, firstly, whether the false positives were reproducible, and secondly, speculating what, exactly, was causing them.

Thus ensued some interesting discussion which ultimately led to a couple of articles from Mark. One at The Conversation and another, slightly more recently, at BBC Future.

I won’t delve into LFT-related science, because Mark has covered it (really, check the video and those articles), but I am going to talk a little bit about pH – the scale chemists use to measure how acidic or alkaline solutions are – because as soon as news of this started to gain traction people, predictably, started trying it out themselves. And that was when things got really interesting.

The buffer included with LFTs is effective at neutralising the pH of solutions, for example, cola

Now, firstly, and importantly: the test kits come with a small vial of buffer solution. Buffers are substances which resist pH changes. As I’ve written before, our bodies naturally contain buffer systems, because keeping the pH of our blood and other body fluids constant is important. In fact, if blood pH varies even a little, you’re in all sorts of serious trouble (which is how we can be certain that so-called “alkaline” diets are a load of hooey). Anyway, the important message is: don’t mix any liquid you’re testing with the contents of that phial, because that will neutralise it.

If you want to try this for yourself, just drop the liquid you want to test directly into the window marked S on the test.

That out of the way, let’s get back to pH. It’s a scale, usually presented as going from 0–14, often associated with particular colours. The 0 end is usually red, the 7 in the middle is usually green, while the 14 end is usually dark blue.

These colours are so pervasive, in fact, that I’ve met more than one person with the idea that acids are red, and alkalis are blue. This isn’t the case, of course. The red/green/blue idea largely comes from universal indicator (UI), which is a mixture of dyes that change colour at different pH values. There’s also a common indicator called litmus (people sometimes mix up UI and litmus, but they’re not the same) which is also red in acid and blue in alkali.

Some species of hydrangea produce pink flowers in alkaline soil, blue in acid soil.

There are actually lots of pH indicators, with a wide variety of colour changes. Phenolphthalein, for example, is bright pink in alkali, and colourless in acids. Bromocresol purple (they have such easy-to-spell names) is yellow in acids, and violet-purple in alkalis.

Many plants contain natural indicators. Just to really mix things up, some species of hydrangea produce flowers that are blue-purple when they’re grown in acidic soil, and pink-red in alkaline conditions.

Bottom line? Despite the ubiquitous diagrams, pH has nothing to do with colour. What it is to do with is concentration. Specifically, the concentration of hydrogen ions (H⁺) in the solution. The more H⁺ ions there are, the more acidic the solution is, and the lower the pH. The fewer there are, the less acidic (and the more alkaline, and higher pH) it is.

In fact, pH is a log scale. When the concentration changes by a factor of 10, the pH changes by one point on the scale.

This means that if you take an acid with pH of 2, and you dilute it 1 part to 10, its pH changes to 3 (i.e. gets one point more alkaline, closer to neutral). Likewise, if you dilute an alkali with a pH of 10 by 1:10, its pH will shift to 9 (again, closer to neutral).

And what this means is that the pH of substances is not a fixed property.

Louder for anyone not paying attention at the back: the pH of substances is not a fixed property!

Yes, we’ve all seen diagrams that show, for example, the pH of lemon juice as 2. This is broadly true for most lemons, give or take, but if you dilute the lemon juice, the pH rises.

Apple juice dropped directly into the test window gives an immediate “positive” result.

I am by no means an expert in commercial, bottled lemon juice, but I reckon a lot of them have water added – which may well explain why @chrismiller_uk was able to get a positive result, while @BrexitClock, using a French bottle of lemon juice, couldn’t.

Mark and I concluded that the pH you need to aim for is probably around 3–4. Go too low, and you don’t get a positive (and you might wipe out the control line, too). Likewise, too high also won’t work.

Myself, I tried apple juice. I couldn’t find the indicator colour key for my indicator paper (I really must clear out the drawers one of these days) but it’s broadly the same as Mark’s cola photo, up above. In other words, the apple juice is about pH 3. And it gives a beautiful positive result, immediately.

One more interesting observation: Mark recorded some time-lapse video comparing orange juice to (sugar-free) cola. It shows the cola test line developing a lot more slowly. We’re not entirely sure why, but it may be pH again: orange juice almost certainly has a lower pH than cola.

For any parents reading this thinking we’re being terribly irresponsible, fear not: as Prof Lorch has made clear in his articles, you can identify a fake by waiting a few minutes and then dropping some of the buffer solution provided in the test window. As I said above, this will neutralise the pH, and the positive test line will disappear. Extra buffer won’t change a genuinely-positive test, because the antibodies bind very tightly (more technical info here). To quote Mark: “you’d need a swimming pool’s worth of buffer to have any chance of washing [the antibodies] off.”

Alternatively, you can just watch your teenager as they do their tests, and make sure they’re not getting up to anything nefarious…

Have you tried to trick an LFT? If you have, share your results! Look us up on Twitter: @chronicleflask and @Mark_Lorch or add a comment below. We’d love to see your photos!

Easy Indicators

Posted on 18 Apr, 2020 by katlday

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:

Oven cleaner, contains sodium hydroxide, ~ pH 13
Plughole sanitiser, contains sodium carbonate peroxyhydrate, ~ pH 10
Baking soda, sodium hydrogen carbonate, ~ pH 8
Washing up liquid, ~ pH 7-8
Tap water, ~ pH 8 (in a hard water area)
Lemon juice, contains citric acid, ~ pH 3
Vinegar, contains ethanoic acid, ~ pH 3
Limescale remover, contains methanoic acid and citric acid, ~ pH 2
Toilet cleaner, contains hydrochloric acid, ~ pH 1

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.

Cleaning chemistry – the awesome power of soap

Posted on 26 Mar, 2020 by katlday

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?

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.

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. COO^–Na⁺ 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. COO^–Na⁺) 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 COO^–Na⁺ stuck on the end. The hydrocarbon end, or “tail”, is hydrophobic (“water-hating”), and only mixes with oily substances. The COO^–Na⁺ end, or “head”, is hydrophilic (“water loving”) and only mixes with watery substances.

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!

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 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!

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

Posted on 26 Jan, 2020 by katlday

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.

Electrolysis Made Easy(ish)

Posted on 19 Dec, 2019 by katlday

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!

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.

Refilling bottles: why it may not be as simple as you thought

Posted on 30 Oct, 2019 by katlday

Two years or so ago most of us had given relatively little thought to single-use plastics. We bought things, we used things, we put the packaging in the bin. Possibly the recycling bin. Hopefully the right recycling bin. And we thought no more about it.

Then Blue Planet II aired on BBC One, specifically episode 7, and suddenly everyone was obsessed with where all this plastic was ending up. Rightly so, since it was clearly ending up in the wrong places, and causing all sorts of havoc in the process.

People started buying reusable cups, eschewing plastic straws and demanding the option of loose fruit and vegetables in supermarkets. Wooden disposable cutlery, oven-cook food containers, and bamboo straws became increasingly common.

And people started to ask more questions about refilling containers. Why do I need a new bottle each time I buy more shampoo or washing up liquid or ketchup, they asked. Why can’t we just refill the bottle? For that matter, couldn’t I take a container to the shop and just… fill it up?

Infinity Foods allow customers to refill containers.

Shops started to offer exactly that. One such place was Infinity Foods, based in Brighton in the UK. Actually, they’d always taken a strong line when it came to recycling and reducing waste, and had been offering refills of some products for years.

Where this gets interesting from a chemistry point of view is a Facebook post they made at the beginning of this month. It said, from the 1st of November, “your empty bottle can only be refilled with the same contents as was originally intended. This includes different brands and fragrances.”

Naturally this spawned lots of comments, many suggesting the change was “daft” and saying things like “I bet it is major corporations not wanting us to reuse the bottle.”

Infinity Foods argued that they were tightening up their policy in order to comply with legislation, specifically the Classification, Labelling and packaging of substances and mixtures (CLP) Regulation (EC) No 1272/2008 and others.

This post, and the comments, got me thinking. I’m old enough, just, to remember the days when random glass bottles were routinely filled with random substances. You wandered into the garage (it was always the garage) and there’d be something pink, or blue, or green, or yellow in a bottle. And it might have a hand-written label, and it might not, and even if it did, the label wasn’t guaranteed to actually be representative of the contents. The “open it and sniff” method of identification was common. The really brave might take their chances with tasting. Home-brew wine might well be next to the lawnmower fuel, and if they got mixed up, well, it probably wouldn’t be fatal.

Probably.

Bottles may be single-use, but they’ve also been designed to be as safe as possible.

You know, I’m not sure we ought to be keen to go back to that, even if it does save plastic. Sealed bottles with hard-to-remove child safety caps, nozzles that only dispense small amounts (making it difficult if not impossible to drink the contents, by accident or otherwise) and accurate ingredients lists are, well, they’re safe.

And we’ve all grown used to them. Which means that now, if I pick up a bottle, I expect the label to tell me what’s in it. I trust the label. If I went to someone else’s house and found a bottle of, say, something that looked like washing up liquid by the sink, I’d assume it was what the label said it was. I wouldn’t even think to check.

You might think, well, so what? You fill a bottle, you know what’s in it. It’s up to you. But what about all the other people that might come into contact with that bottle, having no idea of its origins? What if a visitor has an allergy to a particular ingredient? They look at the label, check it doesn’t contain that ingredient, and use it. Only, someone has refilled that bottle with something else, and maybe that something else does contain the thing they’re allergic to.

Even simpler, someone goes to a shop that sells refills, fills a hair conditioner bottle with fabric softener and doesn’t think to label it. They know what it is, right? They leave it in the kitchen, someone else picks up that bottle, and takes it into the shower. They get it in their eyes and… maybe it causes real harm.

Toilet cleaner must never be mixed with toilet bleach.

Then there are the very real hazards associated with mixing chemicals. One that always worries me is the confusion between toilet cleaner and toilet bleach. Many people have no idea what the difference is. The bottles even look quite similar. But they are not the same substance. Toilet cleaner is usually a strong acid, often hydrochloric acid, while toilet bleach contains sodium hypochlorite, NaClO. Mixing the two is a very bad idea, because the chemical reaction that occurs produces chlorine gas, which is particularly hazardous in a small, enclosed space such as a bathroom.

Okay, fine, toilet bleach and cleaner, noted, check. Is anyone selling those as refills anyway? Probably not. (Seriously, though, if you finish one bottle, make sure you don’t mix them in the toilet bowl as you open the next.)

But it may not be as straightforward as that. Have you ever used a citrus-scented cleaning product? They can be quite acidic. Combine them with bleach and, yep, same problem. What if someone refilled a container that contained traces of a bleach cleaner with one that was acidic, not realising? Not only would it be harmful to them, it could also be harmful for other people around them, including employees, especially if they suffer from a respiratory condition such as asthma.

There are risks associated with the type of container, too. Some plastics aren’t suitable to hold certain substances. Infinity Foods themselves pointed out that some people were trying to find drinking water bottles and plastic milk bottles with cleaning products. These types of bottles are usually made of high-density polyethylene (HDPE). This type of plastic is a good barrier for water, but not oily substances and solvents. Cleaning products could weaken the plastic, resulting in a leak which would be messy at best, dangerous at worst. That’s before we even think about the (un)suitability of the cap.

The type of plastic used to make water bottles isn’t suitable to hold oily substances.

Plus, think of the poor salesperson. How are they supposed to judge, in a shop, whether a particular bottle is safe for a particular product? I wouldn’t feel at all confident about that decision myself. It’s not even always easy to identify which plastic a bottle is made of, and that’s before you even start to consider the potential risks of mixing substances.

In fact, the more you think about it, the more Infinity Foods’ policy makes sense. If you say that you can only refill a bottle with the exact same substance it originally contained, and you insist that the labels have to match, well, that’s easy to check. It’s easy to be sure it’s safe. Yes, it might mean buying a bottle you wouldn’t have otherwise bought, but if you’re going to reuse it, at least it’s just the one bottle.

These concerns all arise from wanting to make sure the world is a safer and healthy place. We do need to cut down on single-use plastics, but taking risks with people’s health to do so surely misses the point.

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.

How are amber teething necklaces supposed to work?

Posted on 30 Sep, 2019 by katlday

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.

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

Posted on 17 Aug, 2018 by katlday

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 ^o C 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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Spectacular Strawberry Science!

Posted on 18 Jun, 2018 by katlday

Garden strawberries

Yay! It’s June! Do you know what that means, Chronicle Flask readers? Football? What do you mean, football? Who cares about that? (I jest – check out this excellent post from Compound Interest).

No, I mean it’s strawberry season in the U.K.! That means there will be much strawberry eating, because the supermarkets are full of very reasonably-priced punnets. There will also be strawberry picking, as we tramp along rows selecting the very juiciest fruits (and eating… well, just a few – it’s part of the fun, right?).

Is there any nicer fruit than these little bundles of red deliciousness? Surely not. (Although I do also appreciate a ripe blackberry.)

And as if their lovely taste weren’t enough, there’s loads of brilliant strawberry science, too!

This is mainly (well, sort of, mostly, some of the time) a chemistry blog, but the botany and history aspects of strawberries are really interesting too. The woodland strawberry (Fragaria vesca) was the first to be cultivated in the early 17th century, although strawberries have of course been around a lot longer than that. The word strawberry is thought to come from ‘streabariye’ – a term used by the Benedictine monk Aelfric in CE 995.

Woodland strawberries

Woodland strawberries, though, are small and round: very different from the large, tapering, fruits we tend to see in shops today (their botanical name is Fragaria × ananassa – the ‘ananassa’ bit meaning pineapple, referring to their sweet scent and flavour.

The strawberries we’re most familiar with were actually bred from two other varieties. That means that modern strawberries are, technically, a genetically modified organism. But no need to worry: practically every plant we eat today is.

Of course, almost everyone’s heard that strawberries are not, strictly, a berry. It’s true; technically strawberries are what’s known as an “aggregate accessory” fruit, which means that they’re formed from the receptacle (the thick bit of the stem where flowers emerge) that holds the ovaries, rather than from the ovaries themselves. But it gets weirder. Those things on the outside that look like seeds? Not seeds. No, each one is actually an ovary, with a seed inside it. Basically strawberries are plant genitalia. There’s something to share with Grandma over a nice cup of tea and a scone.

Anyway, that’s enough botany. Bring on the chemistry! Let’s start with the bright red colour. As with most fruits, that colour comes from anthocyanins – water-soluble molecules which are odourless, moderately astringent, and brightly-coloured. They’re formed from the reaction of, similar-sounding, molecules called anthocyanidins with sugars. The main anthocyanin in strawberries is callistephin, otherwise known as pelargonidin-3-O-glucoside. It’s also found in the skin of certain grapes.

Anthocyanins are fun for chemists because they change colour with pH. It’s these molecules which are behind the famous red-cabbage indicator. Which means, yes, you can make strawberry indicator! I had a go myself, the results are below…

Strawberry juice acts as an indicator: pinky-purplish in an alkaline solution, bright orange in an acid.

As you can see, the strawberry juice is pinky-purplish in the alkaline solution (sodium hydrogen carbonate, aka baking soda, about pH 9), and bright orange in the acid (vinegar, aka acetic acid, about pH 3). Next time you find a couple of mushy strawberries that don’t look so tasty, don’t throw them away – try some kitchen chemistry instead!

Peonidin-3-O-glucoside is the anthocyanin which gives strawberries their red colour. This is the form found at acidic pHs

The reason we see this colour-changing behaviour is that the anthocyanin pigment gains an -OH group at alkaline pHs, and loses it at acidic pHs (as in the diagram here).

This small change is enough to alter the wavelengths of light absorbed by the compound, so we see different colours. The more green light that’s absorbed, the more pink/purple the solution appears. The more blue light that’s absorbed, the more orange/yellow we see.

Interestingly, anthocyanins behave slightly differently to most other pH indicators, which usually acquire a proton (H+) at low pH, and lose one at high pH.

Moving on from colour, what about the famous strawberry smell and flavour? That comes from furaneol, which is sometimes called strawberry furanone or, less romantically, DMHF. It’s the same compound which gives pineapples their scent (hence that whole Latin ananassa thing I mentioned earlier). The concentration of furaneol increases as the strawberry ripens, which is why they smell stronger.

Along with menthol and vanillin, furaneol is one of the most widely-used compounds in the flavour industry. Pure furaneol is added to strawberry-scented beauty products to give them their scent, but only in small amounts – at high concentrations it has a strong caramel-like odour which, I’m told, can actually smell quite unpleasant.

As strawberries ripen their sugar content increases, they get redder, and they produce more scent

As strawberries ripen their sugar content (a mixture of fructose, glucose and sucrose) also changes, increasing from about 5% to 9% by weight. This change is driven by auxin hormones such as indole-3-acetic acid. At the same time, acidity – largely from citric acid – decreases.

Those who’ve been paying attention might be putting a few things together at this point: as the strawberry ripens, it becomes less acidic, which helps to shift its colour from more green-yellow-orange towards those delicious-looking purpleish-reds. It’s also producing more furaneol, making it smell yummy, and its sugar content is increasing, making it lovely and sweet. Why is all this happening? Because the strawberry wants (as much as a plant can want) to be eaten, but only once it’s ripe – because that’s how its seeds get dispersed. Ripening is all about making the fruit more appealing – redder, sweeter, and nicer-smelling – to things that will eat it. Nature’s clever, eh?

There we have it: some spectacular strawberry science! As a final note, as soon as I started writing this I (naturally) found lots of other blogs about strawberries and summer berries in general. They’re all fascinating. If you want to read more, check out…

What’s in your strawberries? (RSC)
The Sweet Chemistry of Five Summer Fruits (Chemical Safety Facts.org)
The Chemistry of Strawberries (Periodic Graphics / Compound Interest)

If you enjoy reading my blog, please consider buying me a coffee (I might spend it on an extra punnet of strawberries, mind you) through Ko-fi using the button below.

the chronicle flask

Tales of interesting chemical tidbits

Tag Archives: acid

Neem: nice, nasty or… not sure?

Faking Lateral Flow Tests: the problem with pH

Easy Indicators

Cleaning chemistry – the awesome power of soap

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

Electrolysis Made Easy(ish)

Refilling bottles: why it may not be as simple as you thought

How are amber teething necklaces supposed to work?

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

Chemists and carbon dioxide

Refreshing fizz

From fertiliser to fizzy drinks

Carbon dioxide calamity?

Don’t despair, DO something…

A new answer from an ancient substance?

Spectacular Strawberry Science!

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Chemists and carbon dioxide

Refreshing fizz

From fertiliser to fizzy drinks

Carbon dioxide calamity?

Don’t despair, DO something…

A new answer from an ancient substance?

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