Continuing with the overenthusiastic children's TV presenter theme (because I feel like I need a lame way to amuse myself today)...
Hey, kids! You know all those batteries that you use to power up all your favourite toys and gadgets? Well, we're going to learn allllllllll about batteries. Well, maybe not all about them, because there's just so much to learn! So, today, we're just going to learn the basics.
The most common kind of battery that you'll see around the shops and in your electronic gizmos is the dry cell. The dry cell is probably called the dry cell because it has a kind of paste, rather than a liquid, inside it.
Here's a picture of the cross-section of a dry cell. The cross-section is what you would see if you cut a dry cell in half, riiiiiight down the middle! But don't try this at home, because it is very very dangerous, and we wouldn't want anyone to get hurt, would we?
Zn(s) à
Zn2+(aq) + 2e-
But wait! If the casing is reacting, how come the casing doesn't corrode? Well, that's because the zinc doesn't just react on its own: it's actually reacting with another substance called magnesium dioxide which is normally found in a powdered form around the graphite rod in the middle of the battery. When they make the batteries, they make sure that there isn't enough magnesium dioxide to react with all of the zinc.
2MnO2(s) + 2NH4+(aq) +
2e- à
Mn2O3(s) + 2NH3(aq) + H2O(l)
After those two reactions happen, the zinc ions from the first reaction combines with the ammonia from the second to produce another ion- a complex ion this time! Ooh, how exciting!
Zn2+(aq) + 4NH3(aq) à [Zn(NH3)4]2+(aq)
Now, eventually, the outer shell of the dry cell does corrode. That's why you sometimes get battery leakages, where the paste inside the battery comes out. This can be very dangerous, so try not to leave your batteries in your electronic devices (or anywhere, really) for so long that they start leaking!
Why does the dry cell corrode even though there isn't enough magnesium dioxide? Well, that paste inside the dry cell doesn't just contain magnesium dioxide- it contains other stuff too, like ammonium chloride and zinc chloride! Ammonium ions in solution can create an acidic environment, which makes the cell corrode after a while.
Since the zinc ions, once formed, are no longer where the original zinc was, dry cells can't be recharged by making the reactions reverse, and you shouldn't try to recharge them either, as that's also quite dangerous! This inability to recharge makes dry cells primary cells, as opposed to secondary cells, which can recharge.
There's also a very similar cell, called the silver oxide cell! These kinds of cells are also primary cells- that is, they can't recharge. They normally appear in button form, like the batteries in watches and in some calculators! They also use the same anode reaction- the conversion of zinc to zinc ions- but silver oxide, rather than magnesium oxide, is consumed at the cathode.
Ag2O(s) + H2O(l)
+ 2e- à
2Ag(s) + 2OH-(aq)
Apparently they've started developing rechargeable versions of these cells, but most silver oxide cells that you'll encounter at this point in time aren't rechargeable. Make sure to only recharge rechargeable batteries, kids!
The next type of cell that I'm going to tell you all about is the lead-acid accumulator! Unlike the dry cell or silver oxide cell, which produce relatively small currents of electricity, the lead-acid accumulator is created to produce lots and lots of current in a very short amount of time. It's often used in cars.
The lead-acid accumulator is actually made up of a bunch of different cells, each with an anode and a cathode. Each electrode has a lead alloy grid, but the anode's grid is packed with finely divided lead and the cathode's grid is divided with lead (IV) oxide powder. You see, when you finely divide something, it makes the reaction proceed much much faster than if you had it all in one single thick lump!
Anode: Pb(s) + SO42-(aq)
à PbSO4(s) +
2e-
Cathode: PbO2(s) + SO42-(aq)
+ 4H+(aq) + 2e- à PbSO4(s) + 2H2O(l)
Also, unlike the previous two batteries, these cells are secondary cells. They can recharge, which makes them much more useful for much longer! You can't recharge them too fast, though, or water might be electrolysed to hydrogen gas and oxygen gas, which would mean that you'd have to top up the battery with distilled water. Not only that, but it might explode and that would defeat the whole idea of the battery being useful for a long time!
2PbSO4(s) + 2H2O(l) à Pb(s) + PbO2(s)
+ 4H+(aq) + 2SO42-(aq)
There are many other kinds of batteries, but there's not enough time to cover them all in today's show. We might be able to talk about them in a future episode. So that's all for now, kids! Stay safe, and look after your batteries!
...
Whew. Now that the show's over, I can go back to being an obnoxious teenager.
By the way, one of my friends' older sisters was watching a whole bunch of kids' shows just to look out for any mistakes that the producers made. Once, she was watching Play School and, at the end, right before the camera was turned off, you could hear someone in the background say, "I hate my job!"
Just some random trivia for you. No, I don't know what episode that is, so I can't help you find it.
Anyway. I'm not going to talk about the other types of batteries, for now at least. Instead, I'm going to talk about corrosion. Yay?
Metals rust and corrode when they're exposed to water and oxygen. How so? you may ask. Well, if you have a look at your data sheet, you'll notice that water and oxygen can be reduced to form hydroxide ions!
O2(g) + H2O(l) + 4e-
à 4OH-(aq)
And, as you probably already know if you've become familiar with the basics of redox, many metals can oxidise. Therefore, what we have here is a redox process of sorts, with the metal oxidising and the oxygen and water reducing.
Bear in mind, though, that those metals right at the bottom of the Standard Reduction Potentials table aren't likely to corrode in the same way. This is because these metals oxidise so readily they tend to react really violently with water to produce hydrogen gas, metal ions and hydroxide ions. Take potassium, for instance:
K(s) + H2O(l) à K+(aq)
+ OH-(aq) + H2(g)
Also, some of the metals near the top of the table won't corrode simply because they're more likely to reduce than water and oxygen are. These metals include gold and silver.
Most other metals, like iron and zinc, will corrode. In fact, iron is very commonly referred to in corrosion questions and examples. My textbook only gives the equations for the rusting of iron. I highly doubt this is really breaching copyright, since you'd probably be able to find all the equations on Google anyway, so here they are:
- Hydroxide ions produced from the reduction of water and oxygen and iron (II) ions produced by the oxidation of iron combine to produce iron (II) hydroxide: Fe2+(aq) + 2OH-(aq) à Fe(OH)2(s)
- Iron (II) hydroxide is easily oxidised (more so than iron) to produce iron (III) oxide: Fe(OH)2(s) + OH-(aq) à Fe(OH)3(s) + e-
- The iron (III) hydroxide is then partially dehydrated to produce hydrated iron oxide, or Fe2O3·xH2O. This is basically just rust, that flaky reddish stuff.
Also, as iron (II) ions are consumed to produce the two hydroxides, the reaction producing these ions is favoured in order to partially restore the number of ions- remember good ol' Châtelier? Since the reaction producing these ions also happens to be the oxidation of iron, the rusting process essentially speeds up the rusting process. Brilliant.
Now, we like our things to last, don't we? So, how can we stop our iron tidbits from corroding? Here are some ways we can prevent a metal from corroding:
- We can coat it with a more reactive metal or connect it via a wire to another more reactive metal. This metal acts as a "sacrificial anode" and will rust instead of the iron or whatever metal that you're trying to protect. If you coat the metal with a less reactive metal, however, then any small scratch on the coating will cause the metal inside to corrode faster. Ditto with linking the metal up to a less reactive metal.
- Cathodic protection: In cathodic protection, the metal to be protected is linked up to the negative terminal of a DC (direct current) power source, while some scrap metal is linked up to the positive terminal. The negative terminal of the power source provides all the electrons that the iron needs to stay as a metal- if any iron dares become an ion, it will simply collect the electrons being fed to it and become iron again.
- Some metals are awesome and actually stop themselves from corroding in the long run. Take aluminium, for example. When aluminium oxidises, it forms aluminium oxide, which actually forms a protective coating that stops any further oxidation. Pretty handy, huh?
That's pretty much all from me on corrosion, unless you feel I haven't covered enough. I might do another post on the different kinds of cells if I can be bothered, but for now I'll just go on to thinking about writing posts about organic chemistry. (Actually writing them might be a little tougher.)
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