Redox in Action: Some Batteries plus Corrosion

Now that we've learned a bit more about redox, let's see how the stuff we've learned applies to the world around us! Yay. Now I sound like an overenthusiastic presenter on a children's TV show.

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?

You know how the different ends of the batteries are labelled positive and negative? Well, they're named that way for a reason. At the anode, which is the negative side of the battery, the zinc casing undergoes a reaction that produces electrons needed for current to flow.

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.

2MnO_2(s) + 2NH₄⁺_(aq) + 2e^- à Mn₂O_3(s) + 2NH_3(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) + 4NH_3(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) + SO₄^2-_(aq) à PbSO_4(s) + 2e^-

Cathode: PbO_2(s) + SO₄^2-_(aq) + 4H⁺_(aq) + 2e^- à PbSO_4(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!

2PbSO_4(s) + 2H₂O_(l) à Pb_(s) + PbO_2(s) + 4H⁺_(aq) + 2SO₄^2-_(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!

O_2(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) + H_2(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:

1. 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)
2. 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^-
3. 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:

1. 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.
2. 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.
3. 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.)

Redox reactions (including the T-word)

The first thing that you have to know is how to write equations for and interpret observations for redox reactions. But we've already done writing equations! you might say, and you're right, since I've already talked about them in my post aptly titled "Redox Equations."  We haven't done observations, though, because observations are totally a whole new ball game. Not.

Basically, the first thing that you need to do is write the equation. Then have a look to see if any of these things happen:

• Gases are bubbled through the liquid
• Gases are produced
• Solids dissolve
• Solids form
• The solution changes colour (use your data sheet to determine the colour of reactant and product ions)

Once you've determined these things, you need to make sure to provide detail. If a gas is bubbled through or if a gas is produced, what colour is the gas, and what does it smell like? Most gases are colourless and odourless, but some have a colour, like chlorine gas which is greenish-yellow. Additionally, most gases are odourless, but there are some that are considered to have a "pungent odour" like hydrogen sulfide (and maybe nitrogen dioxide, I can't remember). If a solid is dissolved or formed, what colour is the solid? (By the way, you have to use what the data sheet tells you. So yes, that means you have to say that copper is "salmon pink.") Also, generally, in a redox reaction, where one solid dissolves and another forms, you can just say that the first solid is coated with the second. I think.

A couple of common ion colours that you should know are the permanganate ion (purple) which can be reduced to manganese (II) ions (pink) and the dichromate ion (orange) which can be reduced to chromate ions (yellow). They're on the data sheet anyway, but they're so common in redox reactions that it's almost worth getting to know them.

What's next? Hm. Stoichiometry. Well, stoichiometry here works the same way as stoichiometry everywhere else, so if you need to brush up on that, head on over to the following posts:

• Simple Calculations Involving Moles
• Reactions and Equations, part deux
• Simple Chemistry Calculations
• Empirical Formulae and Limiting Reagents and Some Other Stuff

Next up: Explain the use of self-indicators in redox titrations. You know how you need an indicator of some kind in acid-base titrations? (If not, go to my post on titrations.) Well, in redox, you generally don't, because common oxidising agents used, like permanganate and dichromate ions, change colour when they're reduced anyway, and it's this colour change that helps you to work out an end point.

Redox titrations are a bit more complex than acid-base titrations because you normally have to go to all the work of sufficiently acidifying the permanganate or dichromate ion solutions otherwise you'll end up with manganese dioxide or chromate ions later on.

Now, a bit more about redox titrations. A primary standard needed here is something that's going to oxidise or reduce, otherwise you wouldn't end up with a redox reaction! Oxalic acid (H₂C₂O₄ , otherwise known as HOOCCOOH which is kinda like Ho-oh with some extra letters added in) is often used for redox titrations as a primary standard because it can be prepared to a high degree of purity and whatnot (my first post on titrations outlines the main characteristics of primary standards).

When using potassium permanganate solution, you have to take all kinds of precautions because it oxidises stuff easily and also decomposes in solution, especially in sunlight. When you prepare the solution, you sometimes have to cover the beaker and boil the solution before filtering it through glass wool into a dark storage bottle. Now, I'm not 100% sure on the science behind all of this, because we haven't really done a lot of redox titrations, but I think the boiling might be to get rid of stuff within the solution that the potassium permanganate might oxidise? I'm not sure exactly what stuff needs to be removed, but there you go. As for the glass wool, I have absolutely no idea whatsoever (and I'm feeling too lazy to look it up right now), so if someone could enlighten me, that would be great! If not, that's okay, you don't have to do all my dirty work for me.

The rest of the titration proceeds in pretty much the same way as an acid-base titration, except for a couple of differences:

1. You have to add sulfuric acid (NOT hydrochloric acid, as the chlorine ions will be oxidised by the permanganate ions) to the solution so that the permanganate or dichromate or whatever solution will form your desired products (i.e. manganese (II) ions rather than manganese dioxide).
2. You have to heat the conical flask before reaction because the reaction proceeds too slowly at room temperature. (Well, too slowly for our purposes, anyway.)
3. You don't have to add indicator because permanganate ions are purple and they should turn pink (or a murky brownish colour if you didn't add enough acid) and dichromate ions are orange and they should turn deep green (or yellow if you didn't add enough acid). However, I've also heard that occasionally the permanganate solution might turn clear instead of pink despite what it says on the data sheet.

Whew. That wasn't so bad. I guess it's easier talking about titrations than actually doing them.

Redox again- Electrolytic cells and some other stuff

(Lack of introduction here because I can't be bothered.)

Previous posts on redox:

• Basics of Redox
• Redox Equations
• Electrolysis
• Applications of Electrolysis
• Ways to Remember Oxidation and Reduction

As with pretty much every other topic we've done this year, the redox stuff this year pretty much takes off from what we did last year. The main difference is that we're going to be talking about electrolytic cells rather than electrolysis- rather than running an electric current through a substance to make a reaction occur, we're using the transfer of electrons in redox reactions to produce electricity. I think.

Anyway, let's take a look at what we need to know.

We still need to know basic stuff like oxidation numbers, as outlined in the Basics of Redox post. Additionally, we need to know how to balance redox equations in normal and in acidic conditions, both of which are outlined in Redox Equations.

By the way, I also know how to balance equations in alkaline conditions. Basically, what you do is you balance the equation as if it's in acidic conditions, then add OH- ions to both sides to turn all H+ into water. Just make sure that water's only left on one side of the equation afterwards.

As an example, let's take the acidic conditions example from the Redox Equations post:

MnO₄^- + 8H⁺ + 5e^- -> Mn²⁺ + 4H₂O

For alkaline conditions, there need to be OH- ions. Accordingly, I add 8OH- to each side (to get rid of H+):

MnO₄^- + 8H⁺ + 8OH^- + 5e^- -> Mn²⁺ + 4H₂O + 8 OH^-

Next I combine the H+ and OH- on the left hand side into water:

MnO₄^- + 8H₂O + 5e^- -> Mn²⁺ + 4H₂O + 8 OH^-

Finally, I cancel out 4 water molecules from each side, and I'm done!

MnO₄^- + 4H₂O + 5e^- -> Mn²⁺ + 8 OH^-

Hmm... what else... "Apply the table of Standard Reduction Potentials to determine the relative strength of oxidising and reducing agents to predict reaction tendency." This isn't too hard. Basically, the higher an element's reduction potential, the more likely it is to be reduced. This, conversely, makes the element a better oxidising agent. Hence, fluorine is a great oxidising agent, and potassium ions are pretty crap in comparison.

When strong oxidants get reduced, they become weak reductants. For example, fluorine, a strong oxidant, can be reduced to fluoride ions which are relatively weak reductants when compared to potassium (which is what you get when potassium ions are reduced).

Another way of thinking about this is by flipping around everything on the Standard Reduction Potentials table to essentially get a Standard Oxidation Potentials table. To get the oxidation potential of the reverse reaction for any reaction on the table, just change the positive or negative sign of the reduction potential.

Okay, I realise that that probably doesn't make sense, so here's an example.

The reduction potential of F₂ + 2e^- à 2F^- is 2.89V. Therefore, the oxidation potential of 2F^- à  F₂ + 2e^- is just -2.89V.

Similarly, if a reaction's reduction potential is negative, the reverse reaction's oxidation potential will be positive.

You might also need to know some of the common oxidants and reductants, though if you're studying Chemistry you'll probably know them already from encountering them often. Possibly the two most common oxidising agents that you'll encounter are the permanganate ion (MnO₄^-) and the dichromate ion (Cr₂O₇^2-), as these are commonly used in redox titrations (titrations AND redox... the two parts of Chemistry that I hate most combined into one... what could be better?).

Other common oxidising agents that you'll need to know, according to my course outline, are oxygen (!), chlorine, the hypochlorite ion (ClO^-- I love Google sometimes), the hydrogen ion, concentrated sulfuric acid and concentrated nitric acid. With regards to the latter two- I think that it's mainly the hydrogen ions in sulfuric acid that oxidise stuff, while that job's given to the nitrate ions in nitric acid, as nitrate ions are better oxidants than hydrogen ions. If you happen to be doing the same Chem course as me you'll notice that nitrate ions have conveniently been left off the data sheet, so here's an equation- two equations, in fact- showing how nitrate ions can be reduced:

NO₃^-_(aq) + 4H⁺_(aq) + 3e^- à NO_(g) + 2H₂O_{(l)     +0.96V}

NO₃^- _(aq) + 2H⁺_(aq) + e^- à NO_2(g) + H₂O_{(l)       +0.80V}

Now, you may ask, out of these two equations, which one will occur in a reaction? Well, judging by the reduction potentials, the top reaction is more likely to happen. I'm assuming that the lower reaction will probably only occur when there aren't enough hydrogen ions. Sorta like how the permanganate ion might reduce to manganese dioxide if the permanganate solution isn't acidified enough.

Due to the nitrate ions being more readily reduced than hydrogen ions in, say, hydrochloric acid, copper will react with concentrated nitric acid, but not with hydrochloric acid. Just some trivia for y'all.

Now let's have a quick look at common reducing agents that you'll need to know. Many metals are good reducing agents, like zinc and magnesium. Hydrogen gas is also a common reducing agent. Iron (II) ions also sometimes oxidise to iron (III) ions (though I'm fairly sure iron itself is a better reductant than its ions), and the chromate ion (C₂O₄^2-) can oxidise to the dichromate ion.

What's up next? Hmm... electrolytic cells! Yay.

Basically, an electrolytic cell is actually made up of two half-cells: one positive, one negative. Each cell consists of some kind of electrode in solution. These electrodes are joined together by a wire which allows for the flow of electrons when the redox reaction between different substances in the cell takes place. I *think* that it's this very flow of electrons that produces electricity. My vague understanding of current from not doing physics and hardly being able to understand the relief teacher/ not really paying attention in year 9 science is that current is basically just the flow of electrons. To complete the circuit, there must be something else connecting these two cells: in the lab, we can just use a salt bridge (a filter paper dipped in solution) with ions that won't react with the ions in either solution. A common solution used for creating a salt bridge is potassium nitrate because neither potassium nor nitrate ions will cause anything to precipitate out of solution.

Enough blabbering on, here's a diagram:

Just like in good ol' electrolysis, the anode is where oxidation takes place and the cathode is where reduction takes place. When stuff gets oxidised at the anode, electrons are removed from whatever is being oxidised and then travel down the wire to the cathode, where they reduce the other substance in question. The salt bridge completes the circuit by allowing charged particles like ions to move between the cells. Since stuff is being oxidised and is becoming more positive in the anode cell, negative ions will migrate towards the anode. Similarly, since stuff is being reduced and is becoming more negative in the cathode cell, positive ions will migrate there.

Also, I need someone to help me out here: is the anode the negative cell, as it produces electrons, while the cathode is the positive cell? This is where I always end up guessing on Chemistry tests. (Well, not always- we've only been asked a total of 2-3 questions about this- 1-2 on the redox test, and one on the exam.)

Sometimes, the electrode in each cell will be a metal and the electrolyte will simply be a solution containing ions of the metal used in the electrode. This, however, is not always the case. The only metal and metal ions that you really need are the ones that are going to take part in the reaction. For example, if you want to utilise a reaction between zinc and nickel (zinc being more likely to oxidise to zinc ions), you do need a zinc electrode in one cell and a nickel ion electrolyte in the other, but you get some degree of choice in the rest. In the zinc cell, the electrolyte can be any solution containing ions (I think that you need ions to help with the flow of electricity- maybe water will suffice in some cases?) that won't react with the zinc. Normally, it's easiest just to use a solution with zinc ions, but I think (I'm not sure) that you can use a solution with manganese or aluminium ions which won't react with the zinc. In the nickel cell, the electrode can be any metal that won't react with the nickel ions- normally graphite or platinum are used as they are inert substances.

What happens if you want to use a gas, like hydrogen gas, as an electrode? Well, that's possible too. What you do is you bubble the gas over a platinum (or carbon, since platinum and carbon are inert) wire/mesh electrode. The wire serves to carry the electrons and the mesh is probably there to increase surface area. I dunno. 

Speaking of hydrogen, there's actually a very special half-cell called the hydrogen half cell which has been assigned a standard reduction potential of 0V. (Did you think that that was an awesome coincidence? I might have. Actually, I can't really remember if I even gave two hoots about it back in Year 11.) All of the other half-cells have had their standard reduction potentials assigned by comparing them to the hydrogen half cell. Here is the hydrogen half-cell, in only some of its glory (well? It's a bit hard for something to maintain its full glory after I've mutilated it on Paint!):

Note that the pressure of the gas, the concentration of the acid and the temperature are all controlled. This is because the standard reduction potentials can change depending on temperature, pressure and concentration. That's right: those Standard Reduction Potentials aren't set in stone. I think most tables show them for 1 mol/L solutions (for substances that are in solutions) at 25 degrees Celsius. (As for gaseous substances, I'm not sure, but I think that they're probably for 101.3kPa, or 1atm. Or 760mmHg if you're that way inclined.)

To work out the standard reduction potentials of other standard cells (cells at 25 degrees C and whatever pressure/concentration is considered "standard"), the other standard cells are hooked up to a standard hydrogen half cell and the voltage recorded. The positive and negative values are assigned depending on what direction the electrons are flowing: if they're going from the hydrogen half cell to the other cell, the values are positive, whereas if they're going from the other cell to the hydrogen half cell, then the values are negative.

Oh, silly me. I've forgotten to tell you all how to calculate the electrical potential of a cell using the table!

Basically, write down the oxidation and reduction half-reactions taking place in the cell (or don't, if you prefer to work mentally). Add the reduction potential for the reduction half-reaction and the oxidation potential for the oxidation half-reaction together to get the electrical potential of the cell. Yay! (I briefly outlined how to work out the oxidation potential earlier in this post.)

By the way, if you get a negative value, you might want to check that you've got your anode and cathode reactions the right way around. Electrical potential values should be positive.

That's pretty much the main stuff covered on redox. Next up will probably be stuff on dry cells (batteries, yay) and corrosion. Oh, and redox titrations, if I can bring myself to write about them.

I still hate redox. And I still hate titrations. They make me really angry for some reason.

Anyway.

TTFN!