Hi. It's Mr. Andersen and this is chemistry essentials video 34 which is on electrochemistry.
In electrochemistry what we're doing is generating current. We're actually generating moving
electrons in a reaction. Or we're requiring those. And so if you don't understand what
a redox reaction is, you may want to look at that video. But a great example of electrochemistry
could be a lemon battery. All we're doing is inserting two pieces of metal into the
lemon and we're going to generate voltage. Now you'd have to hook it up to a pretty sensitive
voltmeter to measure that voltage. But if we hook a number of lemon batteries up in
series we could actually light a little LED. And so electrochemistry is really the study
of redox reactions. And we're separating the reduction and the oxidation portion in space
or time. Now when you do that properly you've created something called an electrochemical
cell. That cell could either be a galvanic cell. It's sometimes also referred to as a
voltaic cell. An example could be a battery. And that's going to generate current. Or we
could have an electrolytic cell. And that's going to require current. An example could
be a rechargeable battery. And so what we're doing remember is separating the reduction
and the oxidation. In review reduction is the gaining of electrons. And oxidation is
the losing of those electrons. And what we're doing is tapping the flow of those electrons.
Now the metal is either a cathode or an anode. And a quick way to remember that is that reduction
always occurs at the cathode. And oxidation and the anode. Or a better way to remember
that is that the consonants R and C are going to pair. And the vowels the O and the A are
going to pair as well. And so what we could do is analyze this cell. We can break it down
into its two half reactions. We could figure out the potential of those two half reactions
and then simply add them together to figure out the standard potential. So how much energy
that cell could produce or require, if it's electrolytic. Now this would be at STP, standard
temperature and pressure. And so if we have variance from that then we could use LeChatelier's
principle to qualitatively measure what's going on. We can also measure stoichiometry
in a cell. So we could use Faraday's laws to qualitative measure that as well. And so
let's start with a simple demonstration. Let's say we take zinc solid and we immerse it in
a copper solution. What will happen over time is we're going to get a reaction. And what
you're going to find is that zinc is going to start to become an ion dissolved in the
solution. And the copper is going to start to form a solid. So you'll actually form this
solid copper on the outside. And so that's a redox reaction. It's reduction. It's oxidation.
There's electrons flowing there but we're not tapping any of those electrons. And so
this would be that whole redox reaction. Again we have the zinc which is losing those electrons
and the copper which is gaining those electrons. And so if we break it down into its 2 half
reactions, it's going to look like this. Oxidation is where the zinc is losing its electrons.
And the copper is gaining those electrons. Those are the two parts. And so in a cell
what we try to do is separate those. But before I actually show you the cell I want to talk
about the reduction potential. In other words that's the potential energy that can really
be released in a reaction like this. And so the zinc, being oxidized is like a waterfall.
In other words it has energy up here and it's releasing that energy. And we could measure
it as 0.76 volts. And so you could look up standard reduction potentials and you'll find
a list of these for a number of different reduction and oxidation reactions. Now the
cool thing is if we run it in the other direction, then it would be negative 0.76 volts. In other
words it would be like trying to move the water up again. It would require that energy.
So if we look at the reduction part this, copper, gaining of electrons, it's going to
have a reduction potential or an E-0 (e-naught) of 0.34 volts. Okay. Now let's actually build
that cell. Let's separate the reduction in the oxidation portion. And so over here we
have the anode. That's going to be where that oxidation takes place. That oxidation of the
zinc. And then we have the cathode. And they're each going to be in a different solution.
So right here we've got copper metal in a copper solution. And then over here we have
zinc ions in a, or excuse me. Zinc solid in a zinc solution. And if we just let it sit
like this nothing will happen. In other words it's not going to generate any voltage if
we had a little voltmeter here up at the top. And the reason why is that we haven't completed
the circuit. And so to do that we're actually going to have to add a salt bridge. A salt
bridge is going to be filled lots of times with a gel. In this case we've got sodium
nitrate. And what it's going to do is it's going to allow the ions to actually flow.
So once we've added the salt bridge, this thing is going to turn on. And electrons are
going to flow from the anode to the cathode. In other words the electrons that are being
lost in the zinc metal, as the zinc becomes a zinc ion, are going to travel through the
wire to the cathode where they're actually going to pick up those electrons. And so let's
talk about what's going on over here in the anode. This is zinc metal, this zinc solid
is actually moving off into an ion in the solution itself. As it does that it's losing
those electrons. And so it's also losing mass. In other words, this anode is going to lose
mass over time. If we look over on the other side, the cathode is actually going to gain
mass. Because that copper that's in solution is going to combine with those electrons that
are coming over from the anode. And it's going to form a copper metal. And so we could see
that the voltage is 1.1 volts. How did I get that 1.1 volts? I simply added up the reduction
potential of those two half reactions. Now how long is this going to occur? It's going
to occur until it reaches some kind of an equilibrium. And so it won't last forever.
Now there's a nernst equation you could use to analyze what happens if it's not really
at standard potential. But that's outside the level of AP Chemistry. What you should
really use is LeChatelier's principle. So what would happen for example if we increased
the amount of zinc on the solution? That would slow this reaction. Or if we were to increase
the amount of copper it's going to speed up this reaction. What happens if we increase
the temperature? That's going to speed up the chemistry. It's going to speed up this
reaction. And it's going to be quicker. In other words we're going to generate more volts
than we would if it was at a cold temperature. And so you should qualitatively understand
how a chemical reaction, or things that would affect a chemical reaction could convert to
things that would affect this cell. Now what's an electrolytic cell? It's really essentially
the opposite of that. In this one what we're going to do is add current. In this case we're
adding a battery on this side. And we're actually going to cause electrons to flow. Great example
of this could be electroplating. When we're actually adding one metal to the other side.
Another example could be rechargeable batteries. And so a rechargeable battery works is you
hook it up to electricity. And what it does is acts as an electrolytic cell. It's going
to store energy inside the cell. Now you unplug it from electricity and now it becomes a galvanic
cell. It can actually release energy. And so we can also use Faraday's laws to really
study stoichiometry. In other words, what's happening to changes in the mass? And there
are two laws. The first one is an equivalence between electricity and the mass. In other
words the more electricity that we apply to a cell, it's directly related to the amount
of mass that a cathode or an anode is going to gain or lose. We also have an equivalence
in Faraday's second law to the mass loss is going to be equivalent to the equivalent weight
of that cathode or anode. What does that mean? Well if we have something like copper and
silver as two of our electrodes, what we would find is that silver, since it has a higher
molecular weight is going to actually lose more mass over time. And so that shouldn't
see surprising if you're looking at it as the losing of atoms from either the anode
of the cathode. And so you should have learned the following. Did you learn to analyze data
from galvanic or electrolytic cells? In other words can you identify what is the reduction
for and what is the oxidation part? And then finally could you apply half reactions.
Again we did that with reduction potential or Faraday's laws to see what's going on in
a cell? Well if you did, you learned what I hoped you would learn. And I hope that was
helpful.