One Flash of Light, One Vision: Carrots, Colour and Chemistry

“White” light is made up of all the colours of the rainbow.

Sometimes you have one of those weeks when the universe seems to be determined to yell at you about a certain thing. That’s happened to me this week, and the shouting has been all about light and vision (earworm, anyone?).

I started the week writing about conjugated molecules and UV spectrometry for one project, was asked a couple of days ago if I’d support a piece of work on indicators for the RSC Twitter Poster Conference that’s happening from 2-3rd March, and then practically fell over a tweet by Dr Adam Rutherford about bacteria that photosynthesise from infrared light in a hydrothermal vent*.

Oh well, who am I to fight the universe?

Light is awesome. The fact that we can detect it is even awesome-er. The fact that we’ve evolved brains clever enough built all sorts of machines to measure other kinds of light that our puny human eyes cannot detect is, frankly, astonishing.

The electromagnetic spectrum covers all the different kinds of light. (Image source)

Let’s start with some basics. You probably met the electromagnetic (EM) spectrum at some point in school. Possibly a particularly enthusiastic physics teacher encouraged you to come up with some sort of mnemonic to help you remember it. Personally I like Rich Men In Vegas Use eXpensive Gadgets, but maybe that’s just me.

The relevant thing here is that the EM spectrum covers all the different wavelengths of light. Visible light, the stuff that’s, well, visible (to our eyes), runs from about 400 to 700 nanometres.

A colour wheel: when light is absorbed, we see the colour opposite the absorbed wavelengths. (Image source)

Now, we need another bit of basic physics (and biology): we see light when it enters our eyes and strikes our retinas. We see colours when only certain wavelengths of light make it into our eyes.

So-called “white” light is made up of all the colours of the rainbow. Take one or more of those colours away, and we see what’s left.

For example, if something looks red, it means that red light made it to our eyes, which in turn means that, somewhere along the way, blue and green were filtered out.

(Before I go any further, there are actually several causes of colour, but I’m about to focus on one in particular. If you really want to know more, there’s this book, although it is a tad expensive…)

Back to chemistry. Certain substances absorb coloured light. We know them as pigments. Carrots are orange, for example, largely because they contain a pigment called beta-carotene (or β-carotene). This stuff appears, to our eyes, as red-orange, and the reason for that is that it absorbs green-blue light, the wavelengths around 400-500 nm.

β-Carotene is a long molecule with lots of C=C double bonds. (Image source.)

Why does it absorb light at all? Well, β-carotene is a really long molecule, with lots of C=C double bonds. These bonds form what’s called a conjugated system. Without getting into the complexities of molecular orbital theory, that means the double bonds alternate along the chain, and they basically overlap and… smoosh into one long thing. (Look, as the saying goes, “all models are wrong, but some are useful,” – it’ll do for now.)

When molecules with conjugated systems are exposed to electromagnetic light, they absorb it. Specifically, they absorb in the ultraviolet region – the wavelengths between about 200 and 400 nanometres. Here’s the thing, though, those wavelengths are right next to the violet end of the visible spectrum – that’s why it’s called ultraviolet after all.

Molecules with really long conjugated systems start to absorb in the coloured light region, as well. And because they’re absorbing violet and blue, possibly a smidge of green, they look… yup! Orangey, drifting into red.

So now you know why carrots are orange. Most brightly coloured fruit, of course, is that way to attract animals and birds to eat it, and thus spread its seeds. As fruit ripens, it usually changes colour, making it stand out better against green foliage and easier to find. This is the link with indicators that I mentioned at the start: many fruits contain anthocyanin pigments, and these often have purple-red colours in neutral-acidic environments, and yellow-green at the more alkaline end. In other words, the colour change is quite literally an indicator of ripeness.

But the bit of the carrot that we usually eat is underground, right? Not particularly easy to spot, and they don’t contain seeds anyway. Why are carrots bright orange?

Modern carrots are mostly orange, but purple and yellow varieties also exist.

Well, they weren’t. The edible roots of wild plants almost certainly started out as white or cream-coloured, as you might expect for something growing underground, but the carrots which were first domesticated and farmed by humans in around 900 CE were, most probably, purple and yellow.

As carrot cultivation became popular, orange roots began to appear in Spain and Germany in the 15th/16th centuries. Very orange carrots, with high levels of β-carotene, appeared from the 16th/17th centuries and were probably first cultivated in the Netherlands. Some have theorised that they were particularly selected for to honour William of Orange, but the evidence for this seems to be a bit slight. Either way, most modern European carrots do descend from a variety that was originally grown in the Dutch town of Hoorn.

In other words, brightly-coloured carrots are a mutation which human plant breeders selected for, probably largely for appearances.

But wait! There was an advantage for humans, too – even if we didn’t realise it straight away. β-carotene (which, by the way, has the E number E160a – many natural substances have E numbers, they’re nothing to be frightened of) is broken up in our intestines to form vitamin A.

Vitamin A is essential for good eye health.

Vitamin A, like most vitamins, is actually a group of compounds, but the important thing is that it’s essential for growth, a healthy immune system and – this is the really clever bit – good vision.

We knew that. Carrots help you see in the dark, right?

Hah. Well. The idea that carrot consumption actually improves eyesight seems to be the result of a World War II propaganda campaign. During the Blitz, the Royal Air Force had (at that time) new, secret radar technology. They didn’t want anyone to know that, of course, so they spread the rumour that British pilots could see exceptionally well in the dark because they ate a lot of carrots, when the truth was that those pilots were actually using radar.

But! It’s not all a lie – there is some truth to it! Our retinas, at the back of our eyes, have two types of light-sensitive cells. Cone cells help us distinguish colours, while rod cells help us detect light in general.

In those rod cells, a molecule called 11-cis-retinal is converted into another molecule called rhodopsin. This is really light-sensitive. When it’s exposed to light it photobleaches (stops being able to fluoresce), but then regenerates. This process takes about thirty minutes, and is a large part of the reason it takes a while for your eyes to “get used to the dark.”

Guess where 11-cis-retinal comes from? Yep! From vitamin A. Which is why one of the symptoms of vitamin A deficiency is night blindness. So although eating loads of carrots won’t give you super-powered night vision, it does help to maintain vision in low light.

Our brain interprets electrical signals as vision.

How do these molecules actually help us to see? Well, when rhodopsin is exposed to light, the molecule changes, which ultimately results in an electrical signal being transmitted along the optic nerve to the brain, which interprets it as vision!

In summary, not only is colour all about molecules, but our whole visual system depends on some clever chemistry. I told you chemistry was cool!

Just gimme fried chicken 😉

*Ah. I sort of ran out of space for the weird hydrothermal bacteria thing. At least one of the relevant molecules seems to be another carotenoid, probably chlorobactene. The really freaking amazing thing is that there seems to be an absorption at 775 nm, which is beyond red visible light and into the infrared region of the EM spectrum. Maybe more on this another day…

If you’re studying chemistry, 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! If you happen to know a chemist, it would make a brilliant stocking-filler! As would a set of chemistry word magnets!

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. If you enjoy reading my blog, and especially if you’re using information you’ve found here to write a piece for which you will be paid, please consider buying me a coffee through Ko-fi using the button below.
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Want something non-sciency to distract you from, well, everything? Why not check out my fiction blog: the fiction phial.


Colour me! STEM Heroes colouring book

Someone reminded me the other day of a podcast I hosted in January 2020, in which I hoped that 2020 would bring everyone lots of good things.

Well, if nothing else, we’ve proved that I definitely don’t have prophetic abilities, eh?

But 2020 hasn’t been all unpleasantness. There have been some bright spots, and I’m about to tell you about one! Back in November the science historian and writer, Dr Kit Chapman (@ChemistryKit), tweeted:

“If I were to commission a colouring book of scientists as heroes/villains (they get to pick what they want to be shown as – superheroes, princesses, wizards etc), would you be up for being a model? Colouring book would be free for all. Just a charity thing for inspiring kids.”

Now, how cool is that idea? Kit set up a GoFundMe which raised (as I write this) over £300, and also sourced twenty different STEM “heroes” to feature in the colouring book. His goal was to ensure multiple ethnicities, gender identities and body types were represented, as well as members of the LGBTQ+ and disabled communities and scientists with mental health disorders. In other words: science is for everyone.

Kit is a science writer (a really good one, read his book) so, of course, he had to include at least one science writer in the book, luckily for me!
 My colouring page is Discworld-themed, because of course it is. It’s based on the Alchemists’ Guild, which on the Disc is… quite an exciting place. To quote a conversation between dwarf Cheery Littlebottom and Sam Vimes in the 19th Discworld book, Feet of Clay:

‘I was quite good at alchemy.’
‘Guild member?’
‘Not any more, sir.’
‘Oh? How did you leave the guild?’
‘Through the roof, sir. But I’m pretty certain I know what I did wrong.’

Like Cheery, I no longer work in a lab, but I do very much enjoy writing about horrible smells, scary acids and everyday chemistry.

You can download a full-size, high-resolution version of my colouring page from here, and you can download the entire book in one go, too — that should keep everyone busy in these slow days between Christmas and New Year!

If you do colour a page — any of them — please come and share it with me: @chronicleflask on Twitter.

I won’t say Happy New Year because, well, that didn’t work out so well last time. So, instead, let’s go with happy end of 2020!

See you all soon and remember, if you’re setting fire to a pudding, do keep it away from the curtains.

If you’re studying chemistry, 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! If you happen to know a chemist, it would make a brilliant stocking-filler! As would a set of chemistry word magnets!

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

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

Monstrous Moles: Happy Mole Day!

Happy Mole Day! It’s the 23rd of October and, at least where I am right now, it’s still between 6:02 am and 6:02 pm, so that means it’s time for chemists to celebrate! Of course, I’m in the U.K., so the date thing doesn’t quite work — for me this is 23/10, not 10/23 — but since there are only 12 months in a year (even in 2020) the British system is a bit unsatisfactory, so I’ll go with the American date format for the day.

There are literally loads of atoms in everything

What’s a mole? Well, to paraphrase Douglas Adams: atoms are small. You just won’t believe how vastly, hugely, mind-bogglingly small they are. I mean, you may think the latest incarnation of the walnut whip is small, but that’s just peanuts to atoms. Or even walnuts.

There are literally loads and loads of atoms in everything. There are so many of the blasted things that the numbers are a real pain to deal with. A teaspoon of table sugar, for example, has about 7,400,000,000,000,000,000,000 sucrose molecules in it, and since each sucrose molecule contains 45 atoms, that’s a whopping 330,000,000,000,000,000,000,000 atoms. And that’s not even a heaped teaspoon.

Even if we used standard form and wrote that last number as, for example, 3.3 x 1023 it’s a bit of a pain. And chemists are far too busy to write things out in full — why do you think they came up with all these symbols in the first place? — so what we do is we pick a convenient amount, which turns out to be 6.022 x 1023, and call that a “mole”. It’s just like calling twelve eggs “a dozen” only, you know, bigger.

I’m not going to explain the origin of the actual number further than this. There’s an awesome graphic here from Compound Interest and, if you want to know more, just click through.

What I am going to do are some… interesting mole calculations. People usually do grains of sand or coins or something. But those are so boooorrring. It’s nearly Halloween, right? I say we go gruesome.

Let’s start with blood!

A healthy adult has about 35 trillion red blood cells in their body at any given moment. (Vampires, presumably, have even more… although… do vampires make their own blood supply? Interesting question…).

35 trillion is a big number, right? A trillion is a million million (on the short scale, which everyone uses, don’t start), 1,000,000,000,000, or 1012, so 35 trillion is 3.5 x 1013.

But that’s only 0.000000000058 of a mole! Even if we count everyone on the planet, we only get to 0.45 of a mole. Yes, that’s right. Even if we drained the blood out of every, single human on the planet, we’d only about half a mole of red blood cells.

Ooh, how about bacteria? We have a lot of those on us, right? In fact, we have more microbes in and on our bodies than human cells! (Well, we can argue about the definition of “human” here, I suppose, but… let’s not.) Apparently there are around 3.8 x 1013 bacteria in our colons which means… damn. This is the blood cells thing all over again, isn’t it? If we took all the humans on the planet, sucked out their gut bacteria (don’t ask) and collected it all together (really, don’t) we’d have, yes, a little under half a mole of microbes.

Don’t tell the tooth fairy’s boss. She really IS scary.

Okay, this is all very well, but it’s not helping us get an idea of scale, is it? All right. Let’s try human teeth. Why not? I mugged the tooth fairy for this one (she’s much tougher than she looks), and it’s about 8 mm long. Adult teeth are a bit larger, of course, but the fairy has less of those. Let’s assume 1 cm to make things easier. That’s 0.01 m. If we had a mole of human teeth they would stack up to… 6.02 x 1021 metres, or 6.02 x 1018 km, or (we need to ramp this up a bit) about 640,000 light years. That would reach a little dwarf galaxy in the constellation of Canes Venatici, somewhere in the general neighbourhood of the Milky Way. Or, alternatively, to Neptune and back…. some 670 million times. Gosh.

What about… hair? A fine human hair is about 0.05 mm across, which means a mole of (fine) hair would be 3.01 x 1019 metres thick. The diameter of the Earth is 12,700,000 metres so that’s about… 2,400,000,000,000 times wider than the Earth. Even Rapunzel might struggle with that much hair.

There are about half a mole of red blood cells in all the humans on the planet.

Hang on, let’s go back to those red blood cells for a minute… okay, if there’s about half a mole of red blood cells in all the humans on the planet, and we assume a single red blood cell is 7.8 μm (0.0000078 m) if we put all those red blood cells in a line it would be…2.3 x 1015 kilometres long. The circumference of the sun is about 4,400,000 kilometres so…

… with all the people on the planet, we could make half a billion rings of blood around the sun, one cell thick.

There’s a thought.

In summary, the mole is a flipping big number. Too big, really, to count anything other than atoms (or molecules, or ions). But it is useful for that.

Happy Mole Day!

Look, these numbers are big, right. I’m going to be amazed if there isn’t some sort of order of magnitude mistake. Just… let me know.

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

Like the Chronicle Flask’s Facebook page for regular updates, or follow @chronicleflask on Twitter. Content is © Kat Day 2020. You may share or link to anything here, but you must reference this site if you do. If you enjoy reading my blog, and especially if you’re using information you’ve found here to write a piece for which you will be paid, please consider buying me a coffee through Ko-fi using the button below.
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Want something non-sciency to distract you from, well, everything? Why not check out my fiction blog: the fiction phial.

More from Genius Lab Gear: Science Word Magnets

Magnets say: the results say we can inhale hot ketones

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

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

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

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

Board shows random magnets

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

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

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

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

The magnets fit with other popular word magnet sets.

Follow this link and the code will be automatically applied.

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

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

Like the Chronicle Flask’s Facebook page for regular updates, or follow @chronicleflask on Twitter. Content is © Kat Day 2020. You may share or link to anything here, but you must reference this site if you do.

Easy Indicators

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

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

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

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

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

This is what you need:

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

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

If you use a plate, cover it with cling film

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

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

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

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

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

Red cabbage indicator with various household substances

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

Here’s a quick summary of what I used:

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

Use the stencils to ensure your hexagons are always perfect!

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

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

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

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

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.

The Chronicles of the Chronicle Flask: 2019

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

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

Newland’s early table of the elements

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

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

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

An atom of Mendeleevium, atomic number 101

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

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

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

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

Britain loves inches.

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

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

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

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

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

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

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

Some STEM Learning trainee teachers, looking very keen!

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

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

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

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

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

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

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

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

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

4OH —> 2H2O + O2 + 4e

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

2H+ + 2e —> H2

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

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

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

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

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

H+ + OH —> H2O

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

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

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

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

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

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

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

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

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

A different version of the experiment…

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

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

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

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

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

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

So, there we are. There’s a lot of potential (haha, sorry) here: you could easily put together multiple class sets of this for a few pounds—the biggest cost is going to be a bulk order of 9V batteries, which you can buy for less than £1 each—and it uses small quantities of innocuous chemicals, so it’s pretty safe. Students could even have their own experiment and not have to work in groups of threes or more, battling with dodgy wires and trippy power-packs (we’ve all been there).

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

Like the Chronicle Flask’s Facebook page for regular updates, or follow @chronicleflask on Twitter. Content is © Kat Day 2019 (photos courtesy of STEM Learning UK and Louise Herbert). You may share or link to anything here, but you must reference this site if you do. If you enjoy reading my blog, please consider buying me a coffee through Ko-fi using the button below.
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Let’s change the way we talk about changes

It’s nearly the end of the school year here in the U.K., traditionally a time for reflecting on what’s gone before and planning ahead for the shiny, new September coming in a mere nine weeks (sorry, teachers!). With that in mind, let’s talk about something that comes up early in most chemistry syllabuses, and which bothers me a little more each time I think about it.

Chemical reactions occur when a match burns.

It’s the concept of chemical and physical changes. For those who aren’t familiar, this is the idea that changes we observe happening to matter fall into two, broad categories: chemical changes, where new substances are made, and physical changes, where no new substances are made.

Examples of chemical changes include things like burning a match, cooking an egg, or the reaction between vinegar and baking soda. Physical changes are largely changes of state, such as melting and boiling, but also include changes such as dissolving salt in water, or grinding limestone chips to powder.

So far, so good. Except… then we start to put descriptors on these things. And that’s when the trouble starts.

multiple choice exam questionThe first problem comes with the idea that “chemical changes are irreversible.” This is often taught in early secondary science as a straight-up fact, and is so pervasive that it’s even appeared in multiple choice exam questions, like the one shown here. The student, for the record, was expected to choose option C, “the change is irreversible.”

Except. Argh. I can tell you exactly why the student has opted for D, “the change is reversible,” and it’s not because they haven’t done their revision. Quite the opposite, in fact. No, it’s because this student has learned about weak acids. And in learning about acids, this student met this symbol, ⇌, which literally indicates a reversible chemical reaction.

Yes, that’s right. Not too long after teaching students that chemical reactions are not reversible, we then explicitly teach them that they are. Indeed, this idea of chemical reversibility is such a common one, such an important concept in chemistry, that we even have a symbol for it.

Now, of course, I can explain this. When we say chemical reactions are irreversible, what we mean is “generally irreversible if they’re carried out in an open system.” In other words, when the wood in that match burns out in the open, the carbon dioxide and water vapour that form will escape to the atmosphere, never to return, and it’s impossible to recover the match to its original state.

The problem is that many chemical reactions occur in closed systems, not least a lot of reactions that happen in solution. Hence, the whole acids thing, where we talk about weak acids “partially dissociating” into ions.

Then there’s that entire topic on the Haber process…

Can I be the only one to think that this is rather a lot of nuance to expect teenagers to keep in their head? It’s nothing short of confusing. Should we really be saying one thing in one part of a course, and the literal opposite in another? To be clear, this isn’t even a GCSE vs. A level thing – these ideas appear in the same syllabus.

Melting is a change of state, in this case from (solid) ice to (liquid) water.

All right, okay, let’s move along to the idea that physical changes are reversible. That’s much more straightforward, isn’t it? If I melt some ice, I can re-freeze it again? If I boil some water, I can condense it back into the same volume of liquid… well… I can if I collect all vapour. If I do it in a closed system. The opposite of the condition we imposed on the chemical reactions. Er. Anyway…

We might just about get away with this, if it weren’t for the grinding bit. If physical changes are truly readily reversible, then we ought to be able to take that powder we made from the limestone lumps and make it back into a nice single piece again, right? Right?

See, this is the problem. What this is really all about is entropy, but that’s a fairly tricky concept and one that’s not coming up until A level chemistry.

Okay. Instead of talking about reversible and irreversible, let’s talk about bond-breaking and bond-forming. That’s fine, isn’t it? In chemical changes, bonds are broken and formed (yep) and in physical changes, they aren’t.


Let’s go back to water for a moment. Water has the formula H2O. It’s made up of molecules where one oxygen atom is chemically bonded to two hydrogen atoms. When we boil water, we don’t break any of those bonds. We don’t form hydrogen and oxygen gas when we boil water; making a hot cup of tea would be a lot more exciting if we did. So we can safely say that boiling water doesn’t involve breaking any bonds, right? We-ell…

Water molecules contain covalent bonds, but the molecules are also joined by (much weaker) hydrogen bonds.

The trouble is that water contains something called hydrogen bonds. We usually do a bit of a fudge here and describe these as “intermolecular forces,” that is, forces of attraction between molecules. This isn’t inaccurate. But the clue is in the name: hydrogen bonds are quite, well, bond-y.

When water boils, hydrogen bonds are disrupted. Although the bonds in individual H2O molecules aren’t broken, the hydrogen bonds are. Which means… bonds are broken. Sort of.

But we’re probably on safe ground if we talk about the formation of new substances. Aren’t we?


What about dissolving? If I dissolve hydrogen chloride gas, HCl, in water, that’s a physical change, right? I haven’t made anything new? Or… have I? I had molecules with a covalent bond between the hydrogen and the chlorine, and now I have… er… hydrochloric acid (note, that’s a completely different link to the one I used back there), made up of H+ and Cl- ions mingled with water molecules.

So… it’s…. a chemical change? But wait. We could (I don’t recommend it) evaporate all that water away, and we’d have gaseous HCl again. It’s reversible.

Solid iodine is silvery-grey, but iodine vapour is a brilliant violet colour.

Hm. What about the signs that a chemical change is occurring? Surely we’re all right there? Fizzing: that’s a sign of a chemical change. Except… are you sure you know the difference between boiling and fizzing? It’s basically all bubbles, after all. Vapour? But, steam is a vapour, isn’t it? Although, on the other hand, water is a product of several chemical reactions. Colour changes? Check out what happens when you heat a small amount of solid, silvery-grey iodine so that it sublimes (spoiler: there’s a colour change).

Is anyone else really confused by now?

You should be. Your students almost certainly are.

There are, in short, more exceptions to every single one of these rules that there are for that “i before e” thing you learned in English (a rule, incidently, which is particularly galling for scientists who constantly have to deal with weights and heights).

Where do we go from here? I think it’s probably time we asked ourselves why we’re even teaching this concept in the first place. Really, it’s there to get students to think about the difference between changes of state and chemical reactions.

I suspect we need to worry about this rather less than we are: most children are very good at identifying changes of state. They do it instinctively. They only start getting confused about it when we teach them a lot of rules which they then try to apply. I’m pretty sure that’s not the way teaching is supposed to work.

A complicated arrangement of chemical glassware

This could definitely be simpler.

If I had my way, I’d ditch the physical and chemical change labels altogether and, instead, just talk about changes of state and chemical reactions. There is precisely one differentiator between these two, and it is: have we made any new stuff? If the answer is no, it’s a change of state. If the answer is yes, then a chemical reaction has occurred. Job done. (And yes, this would squarely define gaseous hydrogen chloride dissolving in water to form hydrochloric acid as a chemical reaction, and I have no problem at all with that.)

I say we change the way we talk about changes: chemistry has a reputation for being tricky, and this sort of confusing, contradictory thing is part of the reason why.

Like the Chronicle Flask’s Facebook page for regular updates, or follow @chronicleflask on Twitter. Content is © Kat Day 2019. You may share or link to anything here, but you must reference this site if you do. If you enjoy reading my blog, please consider buying me a coffee through Ko-fi using the button below.
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A tale of chemistry, biochemistry, physics and astronomy – and shiny, silver balls

A new school term has started here, and for me this year that’s meant more chemistry experiments – hurrah!

Okay, actually round-bottomed flasks

The other day it was time for the famous Tollens’ reaction. For those that don’t know, this involves a mixture of silver nitrate, sodium hydroxide and ammonia (which has to be freshly made every time as it doesn’t keep). Combine this concoction with an aldehyde in a glass container and warm it up a bit and it forms a beautiful silver layer on the glass. Check out my lovely silver balls!

This reaction is handy for chemists because the silver mirror only appears with aldehydes and not with other, similar molecules (such as ketones). It works because aldehydes are readily oxidised or, looking at it the other way round, the silver ions (Ag+) are readily reduced by the aldehyde to form silver metal (Ag) – check out this Compound Interest graphic for a bit more detail.

But this is not just the story of an interesting little experiment for chemists. No, this is a story of chemistry, biochemistry, physics, astronomy, and artisan glass bauble producers. Ready? Let’s get started!

Bernhard Tollens (click for link to image source)

The reaction is named after Bernhard Tollens, a German chemist who was born in the mid-19th century. It’s one of those odd situations where everyone – well, everyone who’s studied A level Chemistry anyway – knows the name, but hardly anyone seems to have any idea who the person was.

Tollens went to school in Hamburg, Germany, and his science teacher was Karl Möbius. No, not the Möbius strip inventor (that was August Möbius): Karl Möbius was a zoologist and a pioneer in the field of ecology. He must have inspired the young Tollens to pursue a scientific career, because after he graduated Tollens first completed an apprenticeship at a pharmacy before going on to study chemistry at Friedrich Wöhler’s laboratory in Göttingen. If Wöhler’s name seems familiar it’s because he was the co-discoverer of  beryllium and silicon – without which the electronics I’m using to write this article probably wouldn’t exist.

After he obtained his PhD Tollens worked at a bronze factory, but it wasn’t long before he left to begin working with none other than Emil Erlenmeyer – yes, he of the Erlenmeyer flask, otherwise known as… the conical flask. (I’ve finally managed to get around to mentioning the piece of glassware from which this blog takes its name!)

It seems though that Tollens had itchy feet, as he didn’t stay with Erlenmeyer for long, either. He worked in Paris and Portugal before eventually returning to Göttingen in 1872 to work on carbohydrates, going on to discover the structures of several sugars.

Table sugar is sucrose, which doesn’t produce a silver mirror with Tollens’ reagent

As readers of this blog will know, the term “sugar” often gets horribly misused by, well, almost everyone. It’s a broad term which very generally refers to carbon-based molecules containing groups of O-H and C=O atoms. Most significant to this story are the sugars called monosaccharides and disaccharides. The two most famous monosaccharides are fructose, or “fruit sugar”, and glucose. On the other hand sucrose, or “table sugar”, is a disaccharide.

All of the monosaccharides will produce a positive result with Tollens’ reagent (even when their structures don’t appear to contain an aldehyde group – this gets a bit complicated but check out this link if you’re interested). However, sucrose does not. Which means that Tollens’ reagent is quick and easy test that can be used to distinguish between glucose and sucrose.

Laboratory Dewar flask with silver mirror surface

And it’s not just useful for identifying sugars. Tollens’ reagent, or a variant of it, can also be used to create a high-quality mirror surface. Until the 1900s, if you wanted to make a mirror you had to apply a thin foil of an alloy – called “tain” – to the back of a piece of glass. It’s difficult to get a really good finish with this method, especially if you’re trying to create a mirror on anything other than a perfectly flat surface. If you wanted a mirrored flask, say to reduce heat radiation, this was tricky. Plus it required quite a lot of silver, which was expensive and made the finished item quite heavy.

Which was why the German chemist Justus von Liebig (yep, the one behind the Liebig condenser) developed a process for depositing a thin layer of pure silver on glass in 1835. After some tweaking and refining this was perfected into a method which bears a lot of resemblance to the Tollens’ reaction: a diamminesilver(I) solution is mixed with glucose and sprayed onto the surface of the glass, where the silver ions are reduced to elemental silver. This process ticked a lot of boxes: not only did it produce a high-quality finish, but it also used such a tiny quantity of silver that it was really cheap.

And it turned out to be useful for more than just laboratory glassware. The German astronomer Carl August von Steinheil and French doctor Leon Foucault soon began to use it to make telescope mirrors: for the first time astronomers had cheap, lightweight mirrors that reflected far more light than their old mirrors had ever done.

People also noticed how pretty the effect was: German artisans began to make Christmas tree decorations by pouring silver nitrate into glass spheres, followed by ammonia and finally a glucose solution – producing beautiful silver baubles which were exported all over the world, including to Britain.

These days, silvering is done by vacuum deposition, which produces an even more perfect surface, but you just can’t beat the magic of watching the inside of a test tube or a flask turning into a beautiful, shiny mirror.

Speaking of which, according to @MaChemGuy on Twitter, this is the perfect, foolproof, silver mirror method:
° Place 5 cm³ 0.1 mol dm⁻³ AgNO₃(aq) in a test tube.
° Add concentrated NH₃ dropwise untill the precipitate dissolves. (About 3 drops.)
° Add a spatula of glucose and dissolve.
° Plunge test-tube into freshly boiled water.

Silver nitrate stains the skin – wear gloves!

One word of warning: be careful with the silver nitrate and wear gloves. Else, like me, you might end up with brown stains on your hands that are still there three days later…

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

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