Genius Lab Gear: The Pocket Chemist

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

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

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

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

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

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

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

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

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

Use the stencils to ensure your hexagons are always perfect!

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

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

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


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.

Electrolysis Made Easy(ish)

Some STEM Learning trainee teachers, looking very keen!

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

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

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

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

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

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

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

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

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

4OH —> 2H2O + O2 + 4e

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

2H+ + 2e —> H2

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

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

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

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

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

H+ + OH —> H2O

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

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

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

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

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

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

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

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

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

A different version of the experiment…

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

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

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

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

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

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

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

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


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

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

Chemical reactions occur when a match burns.

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

Except….

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

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

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

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

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

Except….

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

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

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

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

Is anyone else really confused by now?

You should be. Your students almost certainly are.

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

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

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

A complicated arrangement of chemical glassware

This could definitely be simpler.

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

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


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