There comes a time in every General Chemistry II class when the biochemists-in-training, the pre-med students, and the polymer chemists completely lose interest and the engineers finally understand why chemistry is cool. It is at this time that the chemistry majors stop having an advantage over the biomedical engineering majors and have to, instead, go to them for help. The topic the causes this enormous catastrophe? You guessed it. Electrochemical cells.
Electrochemical cells are these nifty little contraptions through which current flows to make something happen. For voltaic (or galvanic) cells, the current is generated by the cell and used to power things—your cell phone or computer, for example. These are what we refer to as batteries. For electrolytic cells, the current is applied to the cell and forces a non-spontaneous reaction to occur, eventually allowing a person to, for example, bronze their child’s baby booties. (My professor swears it’s a thing.)
Voltaic cells are set up so that there are two compartments in which two half-reactions are occurring. (A half-reaction is either half of a redox reaction, or either half of a reaction where electrons are shuffled around.) Each cell has four very important components—a cathode, an anode, a wire connecting them, and a salt bridge.
The “cat-” component of “cathode” is reminiscent, to most chemistry students, of “cation,” and the two are sort of related. The reduction half-reaction takes place at the cathode; thus, positive cations are always moving toward it. This is how it got its name.
Wait, what is a reduction half-reaction, you ask? Good question. That’s pretty central to this whole concept. You though I was going to skip it, didn’t you?
Let’s say that our electrochemical cell is a Zinc/Copper cell. Because I have a nifty activity series in front of me, I can tell you that copper is going to be the cathode. (I’ll explain that in another post, too.) When something is being reduced, like the copper at the cathode, it is gaining electrons. Therefore, we have this reaction occurring:
Cu2+ (aq) + 2e– → Cu (s)
Or, put in simpler words, copper ions are being taken from the solution the cathode is in, and, with electrons generated at the anode, they are being converted into copper metal.
This is pretty easy to see on paper. It also gives chemistry professors license to ask questions such as, “The copper electrode is increasing in mass. Is it the cathode or the anode?”
You can answer this question now, by the way. It’s the cathode, since metal is being generated and is accumulating on the electrode.
So, if that’s what’s going on at the cathode, what’s happening at the anode? Turns out, since the anode is where oxidation (loss of electrons) takes place, the exact opposite reaction is happening here.
Zn (s) → Zn2+ (aq) + 2e–
See? Here, we have zinc metal from the anode being oxidized into zinc ions (which go into solution) and electrons. That means that the anode is going to lose mass in this case, and give off electrons while it’s doing it.
Now we can see that we’ve got one electrode that uses electrons and one that generates them. So, how do they get from point A to point B, and how does that help power anything?
This is where the wire comes in. When you connect the anode with the cathode, electrons flow from the anode to the cathode by the wire. You can connect that wire to anything your heart fancies, and the electrons will flow through it in the process.
We’ve got electrons flowing from point A to point B via a wire that may or may not be connected to an external circuit. What are we missing?
Well, charges are going to eventually build up at both electrodes. Electrons flowing toward the cathode make it negative, and electrons flowing from the anode make it positive. If these charges are allowed to persist, the electrons will stop moving toward the cathode (the negative electrode), since charges repel each other. Luckily, the compartments of voltaic cells are connected by a salt bridge, a thick gel or filter paper containing ions that can move in and balance out the charges at the electrodes. This completes the circuit, and current can flow continuously.
Relatively simple, aren’t they? Forms of these are found everywhere; they’re dry cell batteries, car batteries, and even hydrogen fuel cells. Of course, electrolytic cells are another game entirely.