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!