Tag Archives: baking soda

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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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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What is Water? The Element that Became a Compound

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November 2018 marks the 235th anniversary of the day when Antoine Lavoisier proved water to be a compound, rather than an element.

I’m a few days late at the time of writing, but November 12th 2018 was the 235th anniversary of an important discovery. It was the day, in 1783, that Antoine Lavoisier formally declared water to be a compound, not an element.

235 years seems like an awfully long time, probably so long ago that no one knew anything very much. Practically still eye of newt, tongue of bat and leeches for everyone, right? Well, not quite. In fact, there was some nifty science and engineering going on at the time. It was the year that Jean-François Pilâtre de Rozier and François Laurent made the first untethered hot air balloon flight, for example. And chemistry was moving on swiftly: lots of elements had been isolated, including oxygen (1771, by Carl Wilhelm Scheele) and hydrogen (officially by Henry Cavendish in 1766, although others had observed it before he did).

Cavendish had reported that hydrogen produced water when it reacted with oxygen (known then as inflammable air and dephlogisticated air, respectively), and others had carried out similar experiments. However, at the time most chemists favoured phlogiston theory (hence the names) and tried to interpret and explain their results accordingly. Phlogiston theory was the idea that anything which burned contained a fire-like element called phlogiston, which was then “lost” when the substance burned and became “dephlogisticated”.

Cavendish, in particular, explained the fact that inflammable air (hydrogen) left droplets of “dew” behind when it burned in “common air” (the stuff in the room) in terms of phlogiston, by suggesting that water was present in each of the two airs before ignition.

Antoine-Laurent Lavoisier proved that water was a compound. (Line engraving by Louis Jean Desire Delaistre, after a design by Julien Leopold Boilly.)

Lavoisier was very much against phlogiston theory. He carried out experiments in closed vessels with enormous precision, going to great lengths to prove that many substances actually became heavier when they burned and not, as phlogiston theory would have it, lighter. In fact, it’s Lavoisier we have to thank for the names “hydrogen” and “oxygen”. Hydrogen is Greek for “water-former”, whilst oxygen means “acid former”.

When, in June 1783, Lavoisier found out about Cavendish’s experiment he immediately reacted oxygen with hydrogen to produce “water in a very pure state” and prove that the mass of the water which formed was equal to the combined masses of the hydrogen and oxygen he started with.

He then went on to decompose water into oxygen and hydrogen by heating a mixture of water and iron filings. The oxygen that formed combined with the iron to form iron oxide, and he collected the hydrogen gas over mercury. Thanks to his careful measurements, Lavoisier was able to demonstrate that the increased mass of the iron filings plus the mass of the collected gas was, again, equal to the mass of the water he had started with.

Water is a compound of hydrogen and oxygen, with the formula H2O.

There were still arguments, of course (there always are), but phlogiston theory was essentially doomed. Water was a compound, made of two elements, and the process of combustion was nothing more mysterious than elements combining in different ways.

As an aside, Scottish chemist Elizabeth Fulhame deserves a mention at this point. Just a few years after Lavoisier she went on to demonstrate through experiment that many oxidation reactions occur only in the presence of water, but the water is regenerated at the end of the reaction. She is credited today as the chemist who invented the concept of catalysis. (Which is a pretty important concept in chemistry, and yet her name never seems to come up…)

Anyway, proving water’s composition becomes a lot simpler when you have a ready supply of electricity. The first scientist to formally demonstrate this was William Nicholson, in 1800. He discovered that when leads from a battery are placed in water, the water breaks up to form hydrogen and oxygen bubbles, which can be collected separately at the submerged ends of the wires. This is the process we now know as electrolysis.

You can easily carry out the electrolysis of water at home.

In fact, this is a really easy (and safe, I promise!) experiment to do yourself, at home. I did it myself, using an empty TicTac box, two drawing pins, a 9V battery and a bit of baking soda (sodium hydrogencarbonate) dissolved in water – you need this because water on its own is a poor conductor.

The drawing pins are pushed through the bottom of the plastic box, the box is filled with the solution, and then it’s balanced on the terminals of the battery. I’ve used some small test tubes here to collect the gases, but you’ll be able to see the bubbles without them.

Bubbles start to appear immediately. I left mine for about an hour and a half, at which point the test tube on the negative terminal (the cathode) was completely full of gas, which produced a very satisfying squeaky pop when I placed it over a flame.

The positive electrode (the anode) ended up completely covered in what I’m pretty sure is a precipitate of iron hydroxide (the drawing pins presumably being plated steel), which meant that very little oxygen was produced after the first couple of minutes. This is why in proper electrolysis experiments inert graphite or, even better, platinum, electrodes are used. If you do that, you’ll get a 1:2 ratio by volume of oxygen to hydrogen, thus proving water’s formula (H2O) as well.

So there we have it: water is a compound, and not an element. And if you’d like to amuse everyone around the Christmas dinner table, you can prove it with a 9V battery and some drawing pins. Just don’t nick the battery out of your little brother’s favourite toy, okay? (Or, if you do, don’t tell him it was my idea.)

Like the Chronicle Flask’s Facebook page for regular updates, or follow @chronicleflask on Twitter. 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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Cooking chemistry: American biscuits

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American biscuitsIt’s come up before of course, but there’s a lot of chemistry in cooking. I do like tinkering with recipes: all that lovely weighing things on digital scales, measuring liquids, working out ratios and tweaking the exact sequence of steps – what more could a chemist want? I spent ages working on my chocolate brownie recipe when I should have been writing up my PhD thesis (it does produce excellent chocolate brownies, so I maintain it was a entirely valid use of my time).

Last week fate transpired to drop more than one reference to ‘biscuits’ in my lap. Now, these were American sources, so I was aware that they weren’t talking about what we call biscuits (and Americans call cookies), not least because in one of them there was talk of making a ‘biscuit sandwich’ that included sausage. Now, I like a chocolate digestive as much as the next person, but I wouldn’t slap a chunk of grilled pork product in between two of them and call it breakfast.

So I decided to try and find a recipe. And, after a bit of faffing around converting Fahrenheit to Celsius and cups to grams (honestly, I do understand the principle of cups and baking by ratio, but is it really easier to measure out a cup of butter than just use scales?) I finally came up with a workable recipe.

Turns out American biscuits are basically sugar-less scones. Who knew.

What’s the chemistry connection? Well, just like scones, the raising agent in American biscuits is baking soda, or sodium hydrogen carbonate. It causes the mixture to rise because it does this when it’s heated:

2NaHCO3 –> CO2 + H2O + Na2CO3

This type of reaction is called thermal decomposition, because the heat is causing the sodium hydrogen carbonate (NaHCO3) to break apart. The carbon dioxide (CO2) is a gas and produces lots of lovely bubbles that make your finished product nice and light.  Water (H2O) is also a product, which helps to keep everything nice and moist.

This clever bit of cookery chemistry starts to happen slowly at 50 oC, but once you get over 200 oC (a more typical baking temperature) it’s pretty fast. So much so that you can bake your biscuits for just 12 minutes or so and they’ll be perfectly risen. Contrary to common belief, there’s no need to add some kind of acid to the mixture (buttermilk is often mentioned). Acids do react with carbonates to produce carbon dioxide, but there’s no need – heat will do the job for you.

Cheese and ham biscuit sandwichSo without further ado, here’s my tinkered recipe. It’s really great this, it literally only takes 10 minutes plus baking time, and you probably have all these ingredients already:

Ingredients

  • 360 g plain white flour
  • 4 tsp of baking soda (sodium hydrogen carbonate)
  • 1 tsp sugar
  • ½ tsp salt
  • 75 g cold (straight from the fridge) unsalted butter, cut into cubes
  • 230 g* milk

(*If you prefer to use a jug this is as close to 230 ml as makes no difference, since milk is mostly water and water has a density of 1 g/ml, but it saves washing up to just stick the bowl on the scales and weigh it.)

Method

  1. Heat the oven to 230 oC.  It needs to be nice and hot, so turn it on in good time.
  2. Measure the dry ingredients in a large bowl and mix them.
  3. Using clean, cold hands rub the butter into the dry ingredients until the mixture resembles fine breadcrumbs and there are no lumps of butter.
  4. Pour in the milk and mix with the flat of a knife until the dough comes together.
  5. Take the dough out of the bowl and place it on a lightly floured surface. Knead it gently a few times until it forms an even ball and has an elastic, ever so slightly sticky, texture.
  6. Press into a rough oblong, about 2 cm thick.  Cut the dough into six roughly equal pieces (you can use cookie cutters, but again, why create unnecessary washing up).  Place these on a greased baking tray.
  7. Bake for about 12 minutes, until the biscuits are golden brown (the colour is, of course, courtesy of another bit of chemistry: the Maillard reaction).
  8. Transfer them to a rack to cool, but no need to leave them too long – they’re best eaten warm!

Split the biscuits in half and fill them with anything you like, savoury or sweet. They’re delicious served plain with lashings of butter. As a more substantial lunchtime snack, try cheese and ham. Lemon curd has also proved a favourite. If you have leftovers they will keep until the next day if wrapped up, and are especially nice toasted and buttered.

And there you are, a metric version of the classic American biscuit recipe, with a bit of chemistry thrown in. I think this might be a first. Heston eat your heart out.

Basic Chemistry

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basic

The other end of the pH scale.

When you start writing a blog it’s hard to predict what people will find most interesting. Inevitably, it’s not what you expected. For example, two of The Chronicle Flask’s most-read posts are about rhubarb and lemons. Perhaps people are more interested in fruit than I ever imagined. Or perhaps I’m getting a lot of hits from people mistakenly looking for recipes.

Or maybe it’s because both feature the ever-interesting topic of acids. In which case, I should probably write something else about acids.

So, this is a post about bases.

Just in case this spectacular bit of contrariness isn’t immediately obvious, bases – some of which are called alkalis (I’m coming to that in a minute) – are at the other end of the pH scale to acids. Acids are the things with a pH value of less than 7, and bases have pH values of more than 7. So basically (hoho), they’re the opposites of acids.

whysoblueI’m using the word base deliberately, and not just because of all the brilliant chemistry puns you can make with it. The more familiar word is probably alkali, but while all alkalis are bases, not all bases are alkalis.

Alkalis are often described as soluble bases. More precisely, alkalis are produced from the metals in group 1 (the ‘alkali’ metals) and group 2 (the ‘alkaline earth’ metals) of the periodic table. The more general term, base, applies to anything that can neutralise an acid. Chemists have another definition: a base is a proton (H+ ion) acceptor, while acids are proton donors (actually chemists have yet another definition, but the proton acceptor one is the one that gets trotted out most often).

The distinction between alkalis and bases does matter to chemists and the two types of substance usually look quite different – bases tend to come in solid lumps or powders (baking soda, for example) and alkalis are more likely to arrive as a solution in a bottle – but in terms of chemistry they both get involved in the same type of chemical reaction, which is neutralising acids.

Indigestion tablet advertWe make use of this all the time, whether we realise it or not. For example if you’re suffering from acid indigestion you probably reach for the indigestion tablets. An advertising campaign for a particular brand of these says that they “turn excess acid into water and other natural substances”. Those ‘natural substances’ are salts – presumably it was decided that the word ‘salt’ had too many negative connotations (which is probably true: how many people would pop a pill that promised to turn into salt in their tummy?) The main ingredient in the tablets in question is calcium carbonate; a base that reacts with stomach acid to produce calcium chloride. Which is definitely a salt, if not the one most people think of when they hear the word.

Tangentially, calcium chloride is also a food additive with the E number E509. It falls into the category of anti-caking agents, which is sort of funny when you think about it.

Anyhoo, that’s one place you use a base (rhyming now as well as punning, sorry). You’re actually making one yourself every time you eat, because your liver produces a substance called bile (bloggers love bile) which helpfully neutralises the acid your stomach produces. If it didn’t, your intestines would get damaged by that acid, so it’s important stuff.

Interestingly, in a lot of the older medical traditions (you know, swallow three leeches with meals, turn around three times under a full moon and bury a toad under a horseradish in a mock turtle) the body’s health depended on the balance of four ‘humors’, or vital fluids: blood, phlegm, ‘yellow bile‘ (choler), and ‘black bile‘. If you had too much of the last two, it was supposed to cause aggression and depression, and in fact the Greek names for them are the root of the words cholera and melancholia.

It’s interesting that in the 21st century many people are obsessed with ‘alkalinizing‘ the body (just check out the comments on that lemons post) when for thousands of years people have understood that too much alkali is probably a bad thing. Public understanding of science has really moved on hasn’t it?

soapBile does something else that’s really quite important in the body, it helps you to digest fats. Bases are generally really good at breaking down fats. This is another thing that’s been known for quite a while, ever since soap was first discovered about (sources vary quite considerably on this) six thousand years ago. Soap is made by a process of saponification, in which fats react with a strong base, usually sodium hydroxide (otherwise known as caustic soda, or sometimes lye). This breaks apart the fat molecules to make glycerol and carboxylate salts (they’re the soap bit). Because of this use, sodium hydroxide features in a famous, and rather gruesome scene, in the film Fight Club.

firediamondNaOH

The fire diamond for NaOH

Because bases are so good at breaking down fats they’re actually surprisingly (or not, if you’ve just watched that Fight Club clip)dangerous, especially because they’re also quite good at breaking down proteins. Your skin is mostly fat and protein, so they can do quite a bit of damage. Remember fire diamonds? The one for sodium hydroxide has a 3 in the blue box, which means that short exposure could cause ‘serious temporary’ or ‘moderate residual’ injury – yikes.

Corrosive hazard symbol

Corrosive hazard symbol

The European hazard symbol is even more alarming, featuring a hand with holes being burned through it. Of course, acids have symbols like these too, but people sort of expect acids to do this kind of stuff. Whereas they’re often (unless they’re chemists) strangely unaware of the dangers of alkalis. For example there’s the a famous, and gruesome, story of the serial killer John George Haigh, who famously dissolved the bodies of his victims in oil drums full of concentrated sulfuric acid. It worked quite well, but he was caught eventually when the police searched his workshop and found sludge containing three human gallstones and part of a denture.

Sulfuric acid is a particularly powerful acid, and is undoubtedly incredibly dangerous stuff, but sodium hydroxide is not much safer. It will cause instantaneous and serious burns, and solid sodium hydroxide gets incredibly hot if it’s added to water. In fact, the water will quickly boil if you’re not careful.

In May last year American Carmen Blandin Tarleton was in the news because she had just received a face transplant. She needed it because her estranged husband had doused her with concentrated sodium hydroxide six years previously. She had undergone fifty-five operations before she made the decision to get the transplant. The pictures are really quite horrific. I won’t reproduce one here; you can see the result of the attack if you follow the link above. Tarleton has also written a book about her experiences. She was left blind and horribly disfigured, with burns to 80% of her body. Doctors described it as “the most horrific injury a human being could suffer”. Sodium hydroxide is not nice stuff.

It’s surprisingly, shockingly, easy to buy sodium hydroxide. Because it’s used in soap-making, you can get it quite easily. It’s even available on Amazon. And of course it’s an ingredient in lots of drain cleaners available in supermarkets. When they say you should wear gloves to handle this stuff, it’s definitely not health and safety gone mad. You really should. Even I would (and I’m really bad about wearing gloves).

So spare a thought for bases. They’re just as interesting, and certainly no nicer or safer than their acidic cousins. In fact, they’re so good at breaking down fat and protein that they could arguably be more dangerous. And next time you’re cleaning out your oven, do remember to wear your gloves.

Are you a chemist and you didn’t know it?

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When I tell people that I’m a chemist, I often get an “oooh, I was really bad at that at school” type response. It’s surprising the number of people that think chemistry has nothing whatsoever to do with their daily lives. Memorably, one acquaintance of an acquaintance (I wouldn’t go so far as to say friend of a friend) once even proclaimed, quite proudly, that the whole of science had nothing to do with her, and she lived her life entirely without it. I was so gobsmacked I didn’t really know where to start, and trust me, that doesn’t happen often.

washing-hands--soap-jpgSo with that in mind, here are five bits of chemistry you do every day. Or at least regularly. You’re a chemist and you didn’t know it!

1. Wash your hands.
Well, we all hope you do this one every day anyway. Soap is very clever stuff. It’s one of the oldest bits of chemistry there is, going back thousands of years, when people first discovered that if they washed their pots with the ashes of cooking fires they got a better result. Soap is made by a process called saponification, where fats are mixed with strong alkalis (traditionally lye: sodium or potassium hydroxide). The fats break apart and form fatty acid salts. What’s clever about those, is that they have a water-loving end (the salt bit) and a water-hating end (the fatty acid bit). So they can grab onto both, and hold the water and oil together. That’s what you do every time you use soap: the dirt ingrained in oil on your skin (nice) can, with the help of those lovely soap molecules, mix with water and so be washed away. Brilliant!

2. Drink a pH indicator.
‘What’ I hear you cry, ‘I do no such thing!’ Ah but do you drink tea (the black kind)? If so, then you do, even if you’ve never noticed. Have you ever put lemon in your tea instead of milk? If not, and you have tea and lemon juice (bottled is fine) in your house, go and try it now. The colour change is really quite lovely to watch. Lemon juice is a source of ascorbic and citric acids, and has a pH of roughly 2-3. You’ll see the same effect with vinegar too, although that mixture wouldn’t be quite so nice to drink. (If you’re feeling adventurous, try some common alkalis such as baking soda or bleach, but DEFINITELY don’t drink those concoctions afterwards…)

3. Carry out combustion.
Ever lit a match? Or a lighter? Started your gas cooker? Turned on your gas boiler? Started your petrol or diesel car? Of course you have. Every single time you do any of those things, the carbon atoms in their molecules are reacting with oxygen to produce carbon dioxide and water. And even if you live under a damp and fireless rock, you’re still doing it – respiration, the process by which all your cells obtain energy – is a form of combustion.

4. Watch some ice float.
Ice floats. Stop press!
We take that for granted, but it’s amazing really. This is a brilliant bit of chemistry that has its tendrils in physics and biology too. Solids don’t generally float on their liquids. Solids are usually more dense than their liquid form, so they sink. But if water behaved like that we wouldn’t have life on this planet, because every time any body of water got really cold it would freeze from the bottom up, taking out all the life swimming in its depths in the process. Since we’re fairly sure that life began in the oceans, evolution would have come to a full stop. But water doesn’t behave like that; water expands when it freezes. Why? Because water has something called hydrogen bonds between its molecules, and as it solidifies these bonds increasingly force the crystalline structure to be very ‘open’. As a result, ice is actually less dense than water, so it floats. This is also why ice is so brilliant at cooling liquids; the warm stuff rises, hits the cold ice and sinks again, creating a sort of cycle called a convection current. Who knew there was so much sciency stuff in your spritzer?

5. Bake a cake.
Food is a rich source of chemistry, just ask Heston. In this case, I’m thinking of baking soda, otherwise known as sodium hydrogencarbonate, or sodium bicarbonate (NaHCO3). When it’s heated above about 70 oC it undergoes a chemical reaction called decomposition. In other words, its molecules break apart without actually needing to react with any other substance. When you put baking soda into your recipe, or use ‘self-raising’ flour (which has it already added), you’re setting it up for this chemical reaction. As the cake cooks, the mixture heats up, and the baking soda does this:
2NaHCO3 –> CO2 + H2O + Na2CO3
The carbon dioxide, CO2, is a gas and it pushes your mixture up and out, causing it to rise. No baking soda chemistry, no lovely, fluffy cake.

So, next time someone tells you they’re rubbish at chemistry, you can point out that they’re doing it every day!