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.