Tag Archives: alkali

Rock bottom: can rocks in your dog’s water bowl protect your lawn?

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fractal image, featuring the hashtag #272sci

Take a look at the Twitter hashtag #272sci

One quick thing before I dive into this month’s post: if you’re a Twitter user, check out my series of very tiny science tweets under the hashtag #272sci. The aim is to explain a science thing in one tweet – without using a thread – and it’s 272 because that’s the number of characters I have to use after including the hashtag and a space. So far I’ve covered leaf colours, frothy milk, caffeine and poisonous millipedes. There will be more to come!

Now, speaking of Twitter, a couple of weeks ago Prof Mark Lorch tweeted about Dog Rocks. Dog… what? I hear you ask (really quite understandably).

Well, it turns out that Dog Rocks are a product that you can buy, and that you put into your dog’s water bowl. Your dog then drinks the water that has been sloshing over the rocks, and, this is where we start to run into trouble, this is meant to have an effect on your dog’s urine. This, in turn, is supposed to protect any grass your dog might then pee on.

photo of a patch of dead grass

Dog urine damages grass

All right, so let’s start somewhere in the vague vicinity of some science: if you have a dog, or even if you’ve just spent some time with someone who has a dog, you’ve probably noticed that dog urine isn’t very kind to grass. Commonly, you see something like the photo here, that is, patches of yellow, dead grass, surrounded by quite luscious green growth.

Why is this? It’s because dog urine – like the urine of all mammals – contains urea, CO(NH2)2. Urea forms in the body when animals metabolise nitrogen-containing compounds, in particular, proteins. It’s essentially a way for the body to get rid of excess nitrogen.

People sometimes confuse urea with ammonia, for reasons that I’ll come to in a moment. But they’re not the same thing. Urea is odourless, forms a pH neutral solution and, if you extract it from the liquid in which it is dissolved, produces solid crystals at room temperature.

Pure ammonia, NH3, by contrast, is a gas at room temperature (boiling point -33.3 ℃), forms alkaline solutions (with pH values greater than 7) and has that pungent ‘ngggh get it away from me!’ smell with which we’re probably all familiar.

Sample pots full of pale yellow liquid

Fresh urine contains urea, but little ammonia

Although these two substances aren’t the same, they are linked: many living things convert ammonia (which is very toxic) to urea (which is much less so) as part of normal metabolism. And it also goes the other way, in a process called urea hydrolysis. This reaction happens in urine once it’s out of the body, too, which is the main reason why, after a little while, urine starts to smell really, really bad.

Okay, fine, but what has this got to do with grass, exactly? Well urea (and ammonia, for that matter) are excellent sources of nitrogen. Plants need nitrogen to grow, but dog urine contains too much, and too much nitrogen is bad – in the same way that too much of pretty much anything nice is bad for humans. It damages the blades of grass and a yellowish dead spot appears, often ringed by some particularly lush grass that, being slightly outside the immediate target zone, caught a whiff of extra nitrogen without being overwhelmed.

Back to Dog Rocks. Interestingly, the website includes an explanation not unlike the one I’ve just given on their fact sheet. What it doesn’t do is satisfactorily explain how Dog Rocks are supposed to change the nitrogen content of your dog’s urine.

photo of a dog drinking water

Dog Rocks are meant to be placed in your dog’s water bowl

The website says that Dog Rocks are “a coherent rock with a mechanically stable framework”. Okay… so… Dog Rocks won’t dissolve or break up in your dog’s water bowl. A good start. It goes on to say, “the rocks provide a stable matrix and a micro-porous medium in which active components are able to act as a water purifying agent through ion exchange” and “Dog Rocks will help purify the water by removing some nitrates, ammonia and harmful trace elements thereby giving your dog a cleaner source of water and lowering the amount of nitrates found in their diet.”

You’ll note they’re using the word nitrate. Nitrates are specifically compounds containing the NO3– ion, but I think they’re using the term in a more general way, to suggest any nitrogen-containing compound (including urea and ammonia). And by the way, nitrates are different from the similar-sounding nitrites, which contain the NO2– ion. Fresh urine from a healthy dog (or human, for that matter) shouldn’t contain nitrite. In fact, a dipstick test for nitrite in urine is commonly used to check for urinary tract infections, because it suggests bacteria are present.

Anyway, nitrates/nitrites aside, it’s the last bit of that claim which really makes no sense. Your dog is not ingesting anything like a significant quantity of nitrogen-containing compounds from its water bowl. Urea comes from the metabolic breakdown of proteins, and they come from your dog’s food.

Photo of puppies eating food that I totally picked because it's cute ;-)

The nitrogen-containing compounds in your dogs’ urine come from their food, not their water

It’s faintly possible, I suppose, that Dog Rocks might somehow filter out some urea/nitrates from urine. But then your dog would have to pee through the Dog Rocks and, honestly, if you can manage to arrange that, you might as well train your dog not to pee on your grass in the first place.

I suggest that there are three possible explanations for the positive testimonials for this product. 1) Owners who use it are inadvertently encouraging their dogs to drink more water, which could be diluting their urine, leading to less grass damage. 2) It’s all a sort of placebo effect: owners imagine it’s going to work, and they see what they’re expecting to see, or 3) they’re all made up.

You decide, but there is absolutely no scientifically-plausible way that putting any kind of rocks in your dog’s water bowl will do anything to stop dog pee damaging your grass. This is £15 you do not need to spend. But hey, you could avoid the money burning a hole in your pocket (see what I did there?) by buying me a coffee… 😉

Check out the Twitter hashtag #272sci here, and support the Great Explanations book project here!

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 – just hit this button:
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Faking Lateral Flow Tests: the problem with pH

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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.

Image

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!

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…
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Onerous ovens: why is cleaning the cooker such a chore?

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As I write Thanksgiving was a few days ago, when most Americans traditionally cook a very large meal based around roasted turkey. Most Brits – and other countries of course – have the same thing coming up soon in the form of Christmas, and there are lots of other celebrations around this time of year that seem to feature cooking and food quite heavily.

Whatever your traditions, then, it’s a time when many of us frown critically at the dark, sticky depths of our oven and wonder if, perhaps, we should attempt to give it a clean. Or at least pay someone else to come and clean it.

Why is oven cleaning such a difficult and unpleasant job, anyway? It’s not that hard to clean other surfaces, is it? Why are ovens so particularly awful?

Well, to explain this, we first need to understand fats.

Fats vaporise during cooking.

Most of the grime in your oven is fat, combined with the carbonised remains of… something or other. The sorts of fats that are common in animal and plant products have boiling points around the 300 oC mark (animal fats typically having higher values than plant oils), but they start to form vapours at much lower temperatures, and certainly at typical cooking temperatures there’s plenty vaporised oil around. Besides, under typical conditions most oils will “smoke” – i.e. start to burn – long before they get close to boiling.

We’re all familiar with the idea that fats don’t mix well with water, and herein lies the problem: all that fatty gloop that’s stuck to the inside of your oven just doesn’t want to come off with standard cleaning methods, particularly when it’s built up over time.

Can chemistry help us here? What are fats, chemically? Well, they’re esters. Which may or may not mean anything to you, depending on how much chemistry you can remember from school. But even if you don’t remember the name, trust me, you know the smell. In particular, nail polishes and nail polish removers contain the simple ester known as ethyl acetate, otherwise known as ethyl ethanoate. (Some people say this chemical smells like pear drops which… only really helps if you know what pear drops smell like. Look, it smells of nail polish, okay?)

Fats are esters (image source)

Anyway, the point is that esters have a particular sequence of atoms that has a carbon bonded to an oxygen, which is bonded to another carbon, which is in turn double-bonded to oxygen. This is a bit of a mouthful, so chemists often write it as COOC. In the diagram here, oxygen atoms are red while carbon atoms are black.

There are actually three ester groups in fat molecules – which explains why fats are also known as triglycerides.

In terms of general chemistry, esters form when a carboxylic acid (a molecule which contains a COOH group) reacts with an alcohol (a molecule that contains an OH group). And this is where it all starts to come together – honest – because you’ve probably heard of fatty acids, right? If nothing else, the words turn up in certain food additive names, in particular E471 mono- and diglycerides of fatty acids, which is really common in lots of foods, from ice cream to bread rolls.

Glycerol is a polyol — a molecule that contains several alcohol groups (image source)

Well, this reaction is reversible, and as a result fats (which are esters, remember) break up into fatty acids and glycerol – which is a polyol, that is, a molecule with several alcohol groups. Or, to look at it the other way around: fats are made by combining fatty acids with glycerol.

And the reason it’s useful to understand all this is that the way you break up esters, and therefore fat, is with alkalis. (Well, you can do it with acid, too, but let’s not worry about that for now.)

Strong alkalis break up fats in a chemical reaction called hydrolysis — the word comes from the Greek for water (hydro) and unbind (lysis) and so literally means “split up with water”. Humans have known about this particular bit of chemistry for a long time, because it’s fundamental to making soap. As I said a few months ago when I was banging on about hand-washing, the ancient Babylonians were making soap some 4800 years ago, by boiling fats with ashes – which works because alkaline compounds of calcium and potassium form when wood is burnt at high temperatures.

The grime in ovens is mostly fat.

The really clever thing about all this is that two things are happening when we mix alkali with fat: not only are we breaking up the fat molecules, but also the substances they break up into are water-soluble (whereas fats, as I said at the start, aren’t). Which makes them much easier to clean away with water. Obviously this is the very point of soap, but it’s also handy when trying to get all that baked-on gunk off your oven walls.

Now, in theory, this means you could get some lye (aka sodium hydroxide, probably), smear it all over your oven and voilà. But I don’t recommend it, for a few reasons. Firstly, it’s going to be difficult to apply, since sodium hydroxide is mostly sold as pellets or flakes (it’s pretty easy to buy, because people use it to make soap).

Sodium hydroxide, sometimes called lye, is often sold in the form of pellets.

But, you say, couldn’t I just dissolve it in water and spray or spread it on? Yes, yes you could. But it gets really, really hot when you mix it with water. So you need to be incredibly careful. Because, and this is my next point, chemically your skin is basically fat and protein, and this reaction we’re trying to do on oven sludge works equally well on your skin. Only, you know, more painfully, and with scarring and stuff. In short, if you’re handing lye, wear good nitrile on vinyl gloves and eye protection.

Actually, regardless of how you’re cleaning your oven you should wear gloves and eye protection, because the chemicals are still designed to break down fats and so… all of the above applies. It’s just that specially-designed oven cleaners tend to come with easier (and safer) ways to apply them. For example, they might come as a gel which you can paint on, and/or with bags that you can put the racks into, and may also be sold with gloves and arm protectors (but rarely goggles – get some separately). They might also have an extra surfactant, such as sodium laureth sulfate, added to help with breaking down grease. The main ingredient is still either potassium hydroxide or sodium hydroxide, though.

Well, possibly, but also not really, if you’re sensible.

As an aside, it makes me smile when I come across an article like this which talks about the “serious” chemicals in oven cleaners and more “natural” ways to clean your oven. The “natural” ways are invariably weak acids or alkalis such as lemon juice or baking soda, respectively. They’re essentially ineffective ways of trying to do exactly the same chemistry.

And okay, sure, the gel and the bag and so on in the modern kits are newer tech, but the strong alkali? Nothing more natural than that. As I said at the start, humans have literally been using it for thousands of years.

A point which really cannot be repeated enough: natural does not mean safe.

Fumes can be irritating to skin, eyes and lungs.

Speaking of which, you will get fumes during oven cleaning. Depending on the exact cleaning mixture involved, these will probably be an alkaline vapour, basically (haha) forming as everything gets hot. Such vapour is potentially irritating to skin, eyes and lungs, but not actually deadly toxic. Not that I recommend you stick your head in your freshly-scrubbed oven and inhale deeply, but you take my point. It might give food a soapy, possibly bitter (contrary to what’s stated in some text books, not all alkalis taste bitter, but do not experiment with this) taste if you really over-do it.

In short, if you’re cleaning your oven yourself: follow the manufacturer’s instructions, make sure your kitchen is well-ventilated, leave the oven door open for a while after you’ve finished and, to be really sure, give all the surfaces an extra wash down with plenty of water.

Put the cleaning off until January – after all, the oven’s only going to get dirty again.

And that’s… it, really. Whether you’re cleaning your own oven or getting someone else to do it for you, the chemistry involved is really, really old. And yes, the chemicals involved are hazardous, but not because they’re not “natural”. Quite the opposite.

Or you could just leave it. I mean, it’s only going to get dirty again when you cook Christmas dinner, right?

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

Easy Indicators

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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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Cleaning chemistry – the awesome power of soap

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

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

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

Soap was first used to clean pots.

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

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

See what I mean about connections?

butyl ethanoate butyl ethanoate

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

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

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

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

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

Still with me? This is where it gets clever. Having mixed our fat with alkali and split our fat molecules up, we have two things: fatty acid salts (hydrocarbon chains with, e.g. 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.

Bars of soap contain sodium stearate.

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

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

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

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

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

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

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

Pure sodium hydroxide is usually supplied as pellets.

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

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

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

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

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

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

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

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

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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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Spectacular Strawberry Science!

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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…

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Alkaline water: if you like it, why not make your own?

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Me* reading the comments section on the Amazing Alkaline Lemons post (*not actually me)

Alkaline water seems to be a trend at the moment. Not quite so much in the UK, yet, but more so in the US where it appears you can buy nicely-packaged bottles with the numbers like 8 and 9.5 printed in large, blue letters on their sides.

It’s rather inexplicable, because drinking slightly alkaline water does literally NOTHING for your health. You have a stomach full of approximately 1 M hydrochloric acid (and some other stuff) which has an acidic pH of somewhere between 1.5 and 3.5. This is entirely natural and normal – it’s there to kill any bugs that might be present in your food.

Chugging expensive water with an alkaline pH of around 9 will neutralise a bit of that stomach acid (bringing the pH closer to a neutral value of 7), and that’s all it will do. A stronger effect could be achieved with an antacid tablet (why isn’t it antiacid? I’ve never understood that) costing around 5p. Either way, the effect is temporary: your stomach wall contains special cells which secrete hydrochloric acid. All you’re doing by drinking or eating alkaline substances is keeping them busy.

(By the way, I’m not recommending popping antacids like sweeties – it could make you ill with something called milk-alkali syndrome, which can lead to kidney failure.)

Recently, a video did the rounds of a woman testing various bottled waters, declaring the ones with slightly acidic pHs to be “trash” and expressing surprise that several brands, including Evian, were pH neutral. The horror. (For anyone unsure, we EXPECT water to have a neutral pH.)

Such tests are ridiculous for lots of reasons, not least because she had tiny amounts of water in little iddy-biddy cups. Who knows how long they’d been sitting around, but if it was any length of time they could well have absorbed some atmospheric carbon dioxide. Carbon dioxide is very soluble, and it forms carbonic acid when it dissolves in water which, yes, would lower the pH.

Anyway, there’s absolutely nothing harmful about drinking water containing traces of acid. It doesn’t mean the water is bad. In fact, if you use an ion exchange filter (as found in, say, Brita filter jugs) it actually replaces calcium ions in the water with hydrogen ions. For any non-chemists reading this: calcium ions are the little sods that cause your kettle to become covered in white scale (I’m simplifying a bit). Hydrogen ions make things acidic. In short, less calcium ions means less descaling, but the slight increase in hydrogen ions means a lower pH.

So, filtered water from such jugs tends to be slightly acidic. Brita don’t advertise this fact heavily, funnily enough, but it’s true. As it happens, I own such a filter, because I live in an area where the water is so hard you can practically use it to write on blackboards. After I bought my third kettle, second coffee machine and bazillionth bottle of descaler, I decided it would be cheaper to use filtered water.

I also have universal indicator strips, because the internet is awesome (when I was a kid you couldn’t, easily, get this stuff without buying a full chemistry set or, ahem, knowing someone who knew someone – now three clicks and it’s yours in under 48 hours).

The pH of water that’s been through a (modern) ion-exchange filter tends to be slightly acidic.

The water in the glass was filtered using my Brita water filter and tested immediately. You can see it has a pH of about 5. The water straight from the tap, for reference, has a pH of about 7 (see the image below, left-hand glass).

The woman in the YouTube video would be throwing her Brita in the trash right now and jumping up and down on it.

So, alkaline water is pretty pointless from a health point of view (and don’t even start on the whole alkaline diet thing) but, what if you LIKE it?

Stranger things have happened. People acquire tastes for things. I’m happy to accept that some people might actually like the taste of water with a slightly alkaline pH. And if that’s you, do you need to spend many pounds/dollars/insert-currency-of-choice-here on expensive bottled water with an alkaline pH?

Even more outlandishly, is it worth spending £1799.00 on an “AlkaViva Vesta H2 Water Ionizer” to produce water with a pH of 9.5? (This gizmo also claims to somehow put “molecular hydrogen” into your water, and I suppose it might, but only very temporarily: unlike carbon dioxide, hydrogen is very insoluble. Also, I’m a bit worried that machine might explode.)

Fear not, I am here to save your pennies! You do not need to buy special bottled water, and you DEFINITELY don’t need a machine costing £1.8k (I mean, really?) No, all you need is a tub of….

… baking soda!

Yep, good old sodium bicarbonate, also known as sodium hydrogencarbonate, bicarb, or NaHCO3. You can buy a 200 g tub for a pound or so, and that will make you litres and litres and litres of alkaline water. Best of all, it’s MADE for baking, so you know it’s food grade and therefore safe to eat (within reason, don’t eat the entire tub in one go).

All you need to do is add about a quarter of a teaspoon of aforementioned baking soda to a large glass of water and stir. It dissolves fairly easily. And that’s it – alkaline water for pennies!

Me* unconvinced by the flavour of alkaline water (*actually me).

Fair warning, if you drink a lot of this it might give you a bit of gas: once the bicarb hits your stomach acid it will react to form carbon dioxide – but it’s unlikely to be worse than drinking a fizzy drink. It also contains sodium, so if you’ve been told to watch your sodium intake, don’t do this.

If I had fewer scruples I’d set up shop selling “dehydrated alkaline water, just add water”.

Sigh. I’ll never be rich.

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Puzzling pool problems?

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We’re half way thorough the Rio 2016 Olypics, and it will have escaped no one’s notice that there have been a few little problems with one of the pools.

Maria Lenk Aquatic Enter, Tuesday, Aug. 9, 2016. (AP Photo/Matt Dunham)

Maria Lenk Aquatic Enter, Tuesday, Aug. 9, 2016. (AP Photo/Matt Dunham)

First, the water turned a mysterious green colour. Then there were reports of a ‘sulfurous’ smell, with German diver Stephan Feck reported as saying it smelled like a “fart”.

The diving pool seemed to be the worst affected, but the water-polo pool next to it also suffered problems, and competitors complained of stinging eyes.

So what on earth was happening? An early suggestion was that copper salts were contaminating the water. It’s not unheard of for copper compounds to get into water supplies, and it would certainly explain the colour; copper chloride solutions in particular are famously greeny-blue. But what about that sulfurous smell? Copper chloride doesn’t smell of sulfur.

Was the strange pool colour due to algae bloom?

Was the strange pool colour due to an algae bloom, like this one in Lake Erie?

The most likely culprit was some sort of algae bloom – in other words rapid algae growth – with the smell probably coming from dimethyl sulfide, or DMS. There’s a singled-celled phytoplankton called Emiliania huxleyi which is particularly famous for producing this smelly compound. In fact, it actually has more than one very important role in nature: the smell is thought to alert marine life that there’s food nearby, but it also seeps into the atmosphere and helps with cloud formation, helping to control our planet’s temperature. Without these reactions, Earth might not be nearly so habitable.

But how did algae manage to grow in the pool? The pool chemicals should have prevented it, so what had happened? An Olympic official then went on to make the comment that “chemistry is not an exact science,” which of course led to much hilarity all around. Chemistry is, after all, incredibly exact. What chemistry student doesn’t remember all those calculations, with answers to three significant figures? The endless balancing of equations? The careful addition of one solution to another, drop by drop? How much more ‘exact’ would you like it to be?

But I had a bit of sympathy with the official, because I suspect that what they actually meant – if not said – was that swimming pool chemistry is not an exact science. And while that, too, is hardly accurate, it is true that swimming pool chemistry is very complicated and things can easily go wrong, particularly when you’re trying to work on an extremely tight schedule. They could hardly, after all, close down all the pools and spend several days carrying out extensive testing in the middle of the sixteen-day-long Olympic Games.

Rio 2016 Olympics Aquatics Stadium (Image: Myrtha Pools)

Rio 2016 Olympics Aquatics Stadium (Image: Myrtha Pools)

When a pool is first built and filled, things are, theoretically, simple. You know exactly how many cubic litres of water there are, and you know exactly how much of each chemical needs to be added to keep the water free of bacteria and other nasties. Those chemicals are added, possibly (particularly in a pool this size) via some kind of automated system, and the pH is carefully monitored to ensure the water is neither too alkaline (basic) nor too acidic.

There’s a certain amount of proprietary variation of swimming pool chemicals, but it essentially all comes down to chlorine, which has been used to make water safe now for over 120 years.

Originally, water was treated to make it alkaline and then chlorine gas itself was added. This produced compounds which killed bacteria, in particular sodium hypochlorite, but the practice was risky. Chlorine gas is extremely nasty stuff – it has, after all, been used as a chemical weapon – and storing it, not to mention actually using it, was a dangerous business.

However, hundreds of people swimming in untreated water is a recipe for catching all kinds of water-borne disease, so it wasn’t long before alternatives were developed.

The Chemistry of Swimming Pools (Image: Compound Interest - click for more info)

The Chemistry of Swimming Pools (Image: Compound Interest – click graphic for more info)

Those alternatives made use of the chemistry that was happening anyway in the water, but  allowed the dangerous bit, with the elemental chlorine, to happen somewhere else. And so hypochlorite salts began to be manufactured to be used in swimming pools.

As the lovely graphic from Compound Interest illustrates, sodium hypochlorite reacts with water to form hypochlorous acid, which in turn goes on to form hypochlorite ions. These two substances sit in an equilibrium, and both are oxidants, which is good because oxidants are good at blasting bacteria. The equilibria in question are affected by pH though, which is one reason why, quite apart from the potential effects on swimmers, it’s so important to manage the pH of pool water.

There are a couple of different chemicals which can be added to adjust pH. Sodium bicarbonate, for example, can be used to nudge the pH up if needed. On the other hand, sodium bisulfate can be used to lower pH if the water becomes too alkaline.

Open-air pools have particular problems

UV light breaks down the chemicals that are used to keep swimming pool water clean.

This can all be managed extremely precisely in an unused, enclosed pool. But once you open that pool up, things become less simple. Open-air pools have a particular problem with UV light. Chlorine compounds are often sensitive to UV – this is why CFCs are such a problem for the ozone layer – and hypochlorite is no exception. In the presence of UV it breaks down in a process called photolysis to form chloride ions and oxygen. This means that outdoor pools require more frequent treatments, or the addition of extra chemicals to stabilise the ‘free available chlorine’ (FAC) levels.

Sadly, I haven’t managed to make it over to Rio, but from what I’ve seen the Aquatic Centre has a roof which opens up, which means that the pool water is indeed being exposed to UV light.

So perhaps the chemical levels simply dropped too low, which allowed algae to proliferate? Possibly aggravated by environmental conditions? Indeed, initially this seemed to be the explanation. FINA, the international governing body of aquatics, issued a statement on Wednesday afternoon which said:

“FINA can confirm that the reason for the unusual water color observed during the Rio diving competitions is that the water tanks ran out of some of the chemicals used in the water treatment process. As a result, the pH level of the water was outside the usual range, causing the discoloration. The FINA Sport Medicine Committee conducted tests on the water quality and concluded that there was no risk to the health and safety of the athletes, and no reason for the competition to be affected.”

This prompted people to wonder how on earth chemical levels were allowed to run out in an event as significant as the Olympics – did someone forget to click send on the order? – but still, it seemed to explain what had happened.

FINA issued a new statement

FINA issued a new statement on Sunday

Until today (Sunday), when more information surfaced as Olympic officials announced that they were going to drain at least one of the swimming pools and refill it. This is no small feat and will involve considerable cost: after all, we’re talking about millions of gallons of water. But it seems to be necessary. As Rio 2016’s director of venue management Gustavo Nascimento said:

“On the day of the Opening Ceremonies of the Games, 80 litres of hydrogen peroxide was put in the water. This creates a reaction to the chlorine which neutralises the ability of the chlorine to kill organics. This is not a problem for the health of anyone.”

Whoops. Yes indeed. Hydrogen peroxide reacts with chlorine to produce oxygen and hydrochloric acid. In fact, hydrogen peroxide is actually used to dechlorinate water which contains levels of chlorine that are too high. It might not be the very worst thing you could add to the water (when you think of all the things that could end up swimming pools) but it’s definitely up there.

Why and how this happened doesn’t, at the moment, appear to be clear. Presumably someone is for the high jump, and not just on the athletics field.

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Are you ok? You look a little flushed.

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PrintYesterday was World Toilet Day (yes, really). This is actually an admirable campaign by WaterAid to raise awareness of the fact that one in three people around the world don’t have access to a safe and private toilet. This, of course, leads to unsanitary conditions which results in the spread of infection and disease. You’ve probably never given it a second thought, but loos literally save lives.

portaloo

Has the TARDIS’ replicator function gone funny?

So, with the topic of toilets in mind, I started thinking about chemical loos. If you live in the UK, the name Portaloo ® will probably spring to mind. This has practically become a generic word for a portable toilet, but it is (like Hoover, Sellotape and others) actually a brand name. I’m told that in America they call them porta-pottys or honey-buckets, which I rather like. In any case, all the chemicals and plastic make them seem like modern inventions, surely?

Actually, not at all. The idea of a self-contained, moveable toilet that you can pick up and take from place to place may be newer, but people have been using chemical toilets for hundreds of years. For example after, ahem, ‘business’ had been completed in an an old-fashioned wooden outhouse – basically a tall box built over a hole in the ground – the user would sprinkle a little lye or lime down the hole to help with the smell.

SodiumHydroxide

Don’t get sodium hydroxide on the toilet seat.

Both of these are strongly basic chemicals. Lye is either sodium hydroxide or potassium hydroxide, and lime is calcium oxide. Both mix with water to form extremely corrosive, alkaline solutions and, incidentally, give out a lot of heat in the process. Both are very damaging to skin. These were the days before health and safety; whatever you did, you had to try not to spill it on the seat.

Urea, a key chemical in urine, reacts with strong alkalis in a process known as alkaline hydrolysis. This produces ammonia, which is pretty stinky (if rather tough on the lungs), so if nothing else that helped to cover up other smells. Ammonia also kills some types of bacteria (which is one reason it’s popular in cleaning products). Flies generally don’t like high concentrations of it either, so that’s another plus.

Alkalis also have another effect in that decomposition of human waste is pH dependent; it works better in acidic conditions. Adding lye or lime raises the pH and slows down this decomposition. On top of this (literally) both lime and lye are hygroscopic: they absorb water. This keeps moisture down and allows a solid ‘crust’ to form on the surface of the waste, making it difficult for any volatile, smelly chemicals to escape. Lovely.

Bleach and ammonia could result in a rocket up your...

Bleach and ammonia could result in a rocket up your…

One word of caution: it’s very, very important you don’t try to clean such an outhouse with any kind of bleach. Bleach, which contains sodium hypochlorite, reacts with ammonia to form hydrogen chloride, chlorine gas and chloramine. None of which are good for your health. Even more dramatically (if this is more dramatic than death – you decide) if there’s lots of ammonia you might get liquid hydrazine, which is used in rocket fuels because it’s explosive. Who knew that toilet chemistry could also be rocket science?

But you don’t find buckets of lye in modern chemical toilets (although, apparently, there are still some people out there using it). So what’s in there? At one time, formaldehyde, otherwise known as methanal, was common. You probably recognise it as embalming fluid; the stuff that Damien Hirst floated that shark in. It’s an extremely effective preservative. Firstly, it kills most bacteria and fungi and destroys viruses, and secondly it causes primary amino groups in proteins to cross-link with other nearby nitrogen atoms, denaturing the proteins and preventing them from breaking down.

shark

Don’t worry, this won’t appear in your chemical toilet.

Interestingly, whilst definitely toxic in high concentrations, formaldehyde is a naturally-occuring chemical. It’s found in the bloodstream of animals, including humans, because it’s involved in normal metabolism. It also appears in fruits and vegetables, notably pears, grapes and shiitake mushrooms. The dose, as they say, makes the poison. I mention this because there are certain campaigners out there who insist it must be completely eliminated from everything, something which is entirely unecessary not to mention probably impossible (just for the hell of it, I’m also going to point out here that an average pear contains considerably more formaldehyde than a dose of vaccine).

All that said, because formaldehyde is extremely toxic in high concentrations, and because it can interfere with the breakdown processes in sewage plants (because it destroys bacteria), formaldehyde isn’t used in toilets so much anymore. In fact, many of the mixtures on sale are explicitly labelled “formaldehyde-free”. Modern formulations are enzyme-based and break down waste by biological activity. They are usually still dyed blue (if you work your way though the colour spectrum, it’s probably the least offensive colour), but usually using food-grade dye. As a result, what’s left afterwards is classed as sewage rather than chemical waste, making it easier to deal with.

Toilet twinning So, this has been brief tour around the fascinating world of toilet chemistry. You’d never have guessed there was so much to it, would you? Now, have you considered twinning your toilet?