Oxidation reduction
Galvanic Cells Redox reactions to drive Galvanic Cells.
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- Let's say I had the reaction where I had some copper ions.
- They have a positive charge of 2.
- It's an aqueous solution.
- There could be some other negative
- ions in that solution.
- Actually we could talk about that in a little bit.
- But we know that when you have these ions in solution, they
- all disassociate.
- So you just have to know you have a bunch of ions just
- floating around in water.
- There could be some other negative ones to neutralize
- the entire solution, but let's just worry about the positive
- ones for now.
- And in that solution, I'm going to put a bar of zinc.
- And what happens is that they essentially switch places.
- The copper wants to jump out of your aqueous solution.
- I should probably draw this.
- Let me draw it.
- This is my solution.
- And in my solution, we make it blue, because it's aqueous, so
- most of it is water.
- I have some copper molecules here in the solution.
- They're a plus-2 charge.
- I could have some other ions that are negative, like some
- nitrate ions or whatever, but they're all disassociated.
- They're all mixing around with the water molecules that are
- polar, and that's why ions disassociate, or salts in
- general, disassociate well in water.
- And I stick a bar of zinc in it.
- Solid zinc. So let me do that.
- So I stick of bar of solid zinc into it.
- And so what this reaction says is that these copper ions, if
- they can get a hold of some electrons and essentially
- become neutral, they would rather go back
- into the solid state.
- And the molecules of zinc out here that are in the solid
- state, if they can lose some electrons, they want to jump
- into the solution.
- So let me write down the whole reaction.
- So if the copper can get some electrons, they're going to go
- into the solid state.
- Copper will go into the solid state.
- And if the zinc can lose some electrons, they'll become
- positive ions or cations.
- And then they will go into the solution, plus-- I'll do the
- same color-- zinc plus 2 plus aqueous.
- Actually just so, a little side note here, these are
- actual ions.
- We've been dealing with oxidation numbers.
- And when you want to know the oxidation number of an ion,
- the oxidation number is the same thing as its charge,
- because it's no longer hypothetical.
- Remember oxidation numbers were hypothetical charges if a
- bond were ionic.
- In this case they are ionic, so the oxidation state, or the
- oxidation number is the same as the charge.
- And just as a side note, the convention, if it's written 2
- plus, this means that this is a real charge.
- If it's written plus 2-- at least this is the way I
- learned, and maybe there's different conventions in
- different places-- then it would be an oxidation state,
- which says, oh, it's not necessarily an ion, but this
- is what it would be if it was an ion.
- So if someone writes plus 2, this is oxidation state.
- If they write 2 plus, they say, hey
- this is a real charge.
- And likewise, if you just have a regular ion that has, like,
- a charge of 1-- so let's say you have Na plus-- this means
- that it's an actual charge.
- If you wanted to write its oxidation state, you would
- write plus 1.
- Small little note, it doesn't probably make a big difference
- in the whole scheme of things, but sometimes it can get
- confusing whether something is an oxidation state or an ion.
- But depending on where you're going to school and what your
- teachers or conventions are, you might want to clarify that
- and just make sure that they use the same conventions that
- at least I remember when I first learned it.
- So anyway, let's go back to this reaction
- that's happening here.
- So let's do the half reactions for this.
- So we start here, and actually we can just look at the
- oxidation states to begin with.
- This guy has an oxidation.
- He has a charge of 2, which is also an oxidation number of 2.
- And he goes to a neutral charge.
- So clearly, he gains two electrons, which means that he
- was reduced.
- He went from a positive charge to a zero
- oxidation state or charge.
- So his charge was reduced.
- Just let me write that.
- So this guy was reduced.
- And let me write his half reaction.
- He essentially started off with some copper in the
- aqueous state, and he got two electrons.
- That's how you get reduced.
- So plus two electrons.
- And he ended up in the solid state of copper.
- And his oxidation number is now 0.
- Now let's see what happened to the zinc. The zinc, it starts
- completely neutral in the solid state.
- No extra electrons or deficiency of electrons.
- But then once the zinc jumps into the solution
- here, he has a charge.
- So he must have lost electrons.
- So the half reaction for zinc-- in this situation you
- have zinc as a solid and he reacts to produce two
- electrons and he's in the water and he's now an ion.
- So let me put that there.
- An aqueous solution.
- So if we did this experiment right here-- and I will tell
- you right now this reaction moves forward.
- This is truly a reaction.
- We're going to talk about, in a second, how badly do the
- electrons essentially want to jump from-- where are they
- jumping from?
- They're jumping from the zinc to the copper.
- So we're going to talk a little bit about how strongly
- they want to jump in the voltage.
- And voltage is really just how badly does charge want to move
- or how much potential energy does charge have. But this
- whole reaction, you just have to think about what's
- happening, is that you have charge moving from-- as the
- zinc jumps into the water, as the zinc jumps into the
- aqueous solution, it's passing on some charge to the copper
- that jumps into the solid form.
- So electrons are moving here.
- And if somehow we could harness these moving
- electrons, we could create a current.
- In this situation, it's all chemical.
- The electrons are moving around, but they're not doing
- it in any organized way.
- Because the solution is all mixed up.
- But this is an interesting situation we're
- dealing with here.
- And we can even talk about how badly the electrons want to
- move in these directions.
- You can look at these half reactions.
- You can look it up in a table, and you can get these energies
- or the voltage is for these reactions.
- And this half reaction right here has a voltage.
- So they'll sometimes right E.
- Sometimes it's an E not right over there, but that's like
- kind of out of the scope of at least a
- first-year chemistry course.
- But they'll say it's 0.34 volts.
- And then this reaction in this direction has a potential
- difference.
- The energy is 0.76 volts.
- If I wrote this in the other way, if I wrote this reaction
- of zinc plus 2 electrons going to zinc solid, then you would
- just switch the sign.
- But what this says is that this state, this state right
- here, is a higher energy than that state.
- So it wants to go in that direction.
- And it wants to do that with an energy per coulomb.
- And you can watch the physics playlist on
- more details of volts.
- But volt units are joules per coulomb.
- This is how badly the electrons
- want to join the copper.
- That's the best way of viewing it.
- And this is a measure of how badly the electrons want to
- leave the zinc. So if you wanted to talk about this
- whole reaction, so this whole reaction that we have up here,
- we could say it's an energy, so the energy of this entire
- reaction up here.
- We know that electrons really want to join the copper with
- this energy, and they really want to leave the zinc with
- this energy.
- So the energy of this entire reaction, how badly do
- electrons want to move from zinc to copper to form this
- over there, is just the sum of the two.
- So you get what's 1.1 volts.
- You're probably saying, hey Sal, all this talk of volts,
- this is interesting and current.
- Is there some way that we could construct a battery or a
- galvanic cell?
- If you don't know what a galvanic cell is, it's what
- we're about to construct right now.
- And the answer is obviously yes, because that's the whole
- point of this video.
- So how can we somehow harness that charge?
- So let's separate the two reactions.
- Let's separate the two.
- So let me put my zinc bar here.
- I'll try to stay true to the colors that I used earlier.
- And let me put my copper bar there.
- Actually maybe I should copy and paste
- what I wrote up here.
- Just so we remember the reaction itself.
- Actually, yeah let me do that.
- Never hurts.
- OK.
- I'll paste it down here.
- Ignore all the squiggly lines around it.
- And let me put but both of my bars in some aqueous solution.
- And there could be some other dissolved ions in there to
- help neutralize the different-- so I put them both
- in an aqueous solution.
- But they're separate.
- And let me label them appropriately.
- So this is my zinc. Zinc right here, a bar of solid sync.
- And this is my bar of solid copper.
- I should do that in a different color.
- So this is copper right here.
- My bar of solid copper.
- And then I have an aqueous solution.
- So what happens if we were to connect
- these two metals, right?
- They're metals.
- If you look at the periodic table they're
- both transition metals.
- They both have the sea of electrons.
- So they both conduct electricity well.
- What happens if you connect these two with a wire?
- So if I connect them with a wire right here like that.
- So we already know that the zinc is just dying to jump
- into the water and lose its electrons, right?
- So let's say that happens.
- So the zinc jumps into the aqueous solution, and it turns
- into a positive ion of zinc. And then its electrons are
- left behind, right?
- So it essentially has two electrons.
- I did the half reaction right here, right?
- The zinc jumps into the solution and it leaves behind
- two electrons right here.
- Two electrons here, they really want to jump off of the
- zinc.
- Now I have a lot of copper out of here, out here in the
- solution, that really want to jump on to this bar of copper.
- But right now they're all ions.
- These are all plus 2 ions.
- And they really want to jump on to the bar and join all of
- these other copper ions here.
- But in order to jump on to that, they
- have to get some electrons.
- Each of these guys have to grab two electrons.
- And we know that kind of potential, or the
- electromotive force with which they want to grab those
- electrons is 0.34 volts.
- These guys want to jump here and grab electrons from
- someplace with plus 0.34 volts.
- These guys want to jump off of the bar and leave their
- electrons behind with 0.76 volts.
- So what do you think's going to happen if we have a
- conducting wire here?
- Well those electrons are, once this zinc just jumps off,
- those electrons are going to travel through the wire.
- And then they're going to be available, because these are
- surplus electrons, to merge with these copper ions that
- want to jump back on to the bar.
- So these copper ions are going to jump back on to the bar.
- But from this reaction what's happening?
- You're having a current flow.
- Current is actually flowing.
- And what is the voltage across this?
- How badly does an electron want to go from the zinc on
- this bar to over here, join with some copper in the
- solution and let that copper join on to that bar?
- It wants to do it with plus 1.1 volts.
- Now the question is the electrons are flowing in that
- way, what is the cathode and what is the anode?
- So the cathode is the positive terminal of, well, you can
- already imagine, this is a battery of sorts, right?
- Because we are able to generate a current by having a
- voltage difference across two terminals.
- The cathode is the positive terminal.
- That's where the electrons want to go to.
- And then the anode-- and you can familiarize yourself with,
- you know, just a cation is a positive ion and anion is a
- negative ion.
- So the anode is a negative terminal.
- So what's the negative terminal here?
- The negative terminal is where the electrons come from.
- See, electrons are coming from the zinc jumping off.
- So is your negative terminal.
- And what's the positive terminal?
- That's where the electrons want to go to.
- They want to go down here and be available for these coppers
- that want to jump out of the solution.
- So this is our positive terminal, or our cathode.
- This is the cathode.
- This is the anode.
- And if someone says, hey, does oxidation or reduction occur
- at the cathode or the anode?
- Well what's occurring right here?
- We have zinc in its solid state right here.
- We have solid zinc jumping in-- and it's neutral-- and it
- jumps into the solution and it leaves two electrons behind.
- So it's losing electrons.
- When you lose electrons you're being oxidized.
- So this is where the oxidation is occurring.
- Likewise, on this side you have copper ions that are
- already in the solution, and they're gaining electrons.
- So their charge is being reduced.
- So the cathode is where reduction is taking place.
- Reduction.
- Now you're probably saying, oh Sal, this is a nice cute
- little, you know, battery-- or this is called a galvanic
- cell-- that, you know, you've created this current.
- But I see a problem here.
- Because as this happens, you're going to have more and
- more positive zinc ions jump into the solution here.
- You have more and more positive zinc ions in here as
- this progresses.
- And you're going to have fewer and fewer copper ions in here.
- So if you assume that there were some other, let's say
- there was, I mean, maybe to keep the solution neutral--
- the copper is positive-- maybe you had some chloride in
- there, chloride anions that were keeping
- this solution neutral.
- And maybe you had some chlorine anions here that were
- keeping, well, it started negative but then it maybe
- gets a little bit more neutral as the zinc jumps in.
- But once this happens a lot then it becomes
- very positive, right?
- This is becoming more and more positive as more and
- more zinc jumps in.
- And this is becoming more and more negative as more copper
- jumps on to this bar.
- So what can we do to solve that problem?
- What they do is-- and this is just one way to solve the
- problem, there's actually many but this probably what you
- might encounter in your first-year chemistry class--
- is they create something called a salt bridge.
- And what they do is they create a, I guess we could
- call it a pseudo-aqueous solution of salt.
- Maybe it's potassium chloride.
- So in here we have potassium.
- Potassium has a plus charge.
- It's aqueous, plus chlorine with a
- negative charge aqueous.
- And I say pseudo-aqueous because it's not really
- super-liquid water.
- What they do is they try to make it more viscous so it
- just doesn't mix completely.
- If this was pure water, then all of it would just fall into
- the solution, and you would really have nothing up here.
- It would be a useless device.
- But what they do is they create a gel that's still
- aqueous enough that the ions can move around.
- The sodium and chloride ions can move around.
- So let's say that those are the-- sorry, not the sodium,
- the potassium-- let's say that's the potassium.
- And let's say the chlorine ions are-- I already did it in
- that blue-- so chlorine are the minus.
- And what happens is as we have this oxidation reduction
- reaction that's causing these electrons to travel from the
- anode to the cathode, this becomes positives, this
- becomes negative.
- The salt bridge allows the negative components of the
- ionized salt and just waste the chlorine to go into the
- solution to neutralize it.
- So they'll end up having more and more chlorine as
- the zinc pops off.
- If you were to evaporate the solution you'd end up with a
- bunch of zinc chloride.
- And on this side, to keep the balance here, you'll have a
- lot of the positive parts, components, of the salt.
- So the sodium would jump in to replace the copper that's
- jumping out.
- So when you have this whole thing it enables the current
- to keep flowing without having a bad effect on the charge.
- Because, you know, what happens if this
- becomes really positive?
- If this becomes really positive-- I'll do it in
- vibrant color because there's no salt bridge-- at some point
- the electrons are going to say, hey, I can either go over
- here and drive this reaction or, hey, I got a lot of
- positive stuff hanging out here.
- Maybe I don't want to leave it all.
- So you have to do something to neutralize the charge here.
- Likewise, if this becomes really negative from all the
- chlorine atoms jumping out, electrons aren't going to want
- to go here.
- Because like charges repel each other.
- So that's what having the salt bridge neutralizing the
- solution does for you.
- It keeps it from getting negative, and then our
- galvanic cell can keep operating.
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At 5:31, how is the moon large enough to block the sun? Isn't the sun way larger?
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