Oxidation reduction
Introduction to Oxidation States Oxidation and reduction. Oxidation states.
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- Let's have the molecule sodium chloride.
- Let's look at the periodic table.
- Sodium, it's in alkali metal.
- It has one valence electron.
- It's in Group 1.
- And so when it only has one, it really
- wants to give it away.
- And chlorine is a halogen, and it's only one away.
- It just needs one electron to have the full eight valence
- electrons in its outermost shell.
- So it really wants to take 1 electron.
- And we've talked about this story multiple times,
- especially it early in the chemistry playlist. So what we
- know what's going to happen.
- Sodium, since it wants to lose an electron, is going to lose
- its electron.
- And it's going to give it to chlorine.
- And chlorine is going to gain it.
- And so you're going to have a situation where sodium is
- going to have a plus 1 charge, because it lost an electron.
- And then chlorine will have a minus 1 charge because it
- gained an electron.
- And then they'll be attracted to each other because this one
- is now a positive atom and that one is
- now a negative atom.
- And just the Coulomb force will make them want to be with
- each other.
- And we call this, of course, an ionic bond.
- I haven't taught you anything you don't know yet.
- And in an ionic bond, you literally had a loss of
- electron from one compound to another.
- Now.
- that's a very cut-and-dry situation.
- But we saw other situations where wasn't quite as
- cut-and-dry.
- For example, H2O.
- Where you have a oxygen-- but you've seen this multiple
- times already, if you're watching
- this playlist in order.
- And it's bonded to 2 hydrogens.
- But we know that oxygen is much, much more-- in fact,
- nitrogen, oxygen, and fluorine are the most electronegative
- atoms, which means that they like to hog electrons the
- most. So when you have oxygen bonded with pretty much
- anything-- but in particular in this case, hydrogen-- it's
- going to to pull most of the electrons in its direction.
- So even though these are covalent bonds, the hydrogen
- is still sharing the electrons that it
- has bonded with oxygen.
- The electrons spend most of their time on the oxygen side
- of this whole affair.
- So you have a partial negative charge on the oxygen end of
- the molecule.
- And you have partial positive charges on the hydrogen.
- And we talked about this.
- And this is what leads to hydrogen
- bonding, and all of that.
- What we're going to do here is introduce-- it's almost an
- intellectual tool.
- We know that this isn't an ionic bond.
- We know that this is actually a covalent bond, that this
- atom is shared.
- But it spends most of its time at the more
- electronegative atom.
- That's why you have this partial negative charge.
- So we create this convention called oxidation states.
- Let me write that down.
- And I'll clarify what the word means in a second.
- And this is essentially assigning-- With an ionic
- compound, you naturally assign a charge, because each atom
- really does have a charge.
- But let's say we want to live in a world where we don't like
- this partial charge, partial negative.
- We want to say, look, if this were, hypothetically, an ionic
- bond, what would it look like?
- Who would gain the electron and who
- would lose the electron?
- So in the case of water, if you were forced to say OK, who
- gained the electron?
- You'd say oxygen gained 2 electrons from the hydrogens,
- and the hydrogens lost each 1 electron to the oxygen.
- So the hydrogen would have a plus 1 charge, each of them.
- And then the oxygen would have a minus 2 charge.
- Now, I want to be very clear with this.
- This isn't what really happened.
- This is just our little intellectual game that we're
- playing called an oxidation state.
- It's going to be really useful later to understand why some
- reactions occur.
- But I just want to be very clear.
- This is a hypothetical charge if these were ionic bonds.
- So you're just saying, whoever is the more electronegative--
- and remember, electronegativity, it goes
- from the bottom left to the top right.
- So these are the most electronegative atoms, which
- means they love to hog electrons the most. These are
- the least. Or you could also call them the most
- electropositive, which means they like to give away
- electrons the most.
- So what you would do is you say, OK.
- The more electronegative, let's just say that they
- actually get the electrons.
- And that the more electropositive atoms, they
- give the electrons.
- Even though we know that it's something in between.
- It's actually partial.
- Now these numbers.
- These hypothetical ionization of these hydrogen molecules.
- This is called their oxidation number.
- The oxidation number of hydrogen in H2O is plus 1.
- So we could write a plus 1 for each of the hydrogens there.
- And then oxidation number for the oxygen is minus 2.
- And the way we talk about it, we say that the hydrogen
- molecules here have been oxidized.
- And it's kind of like saying that you've been-- In this
- case, they were oxidized by oxygen.
- And this is actually a very confusing point, or at least
- it was to me initially.
- Because when I first learned about oxidation, I was like,
- oh, that's what oxygen does to things.
- Because the word has the word, oxygen, or at least the
- beginning part of oxygen.
- So I thought, oh, oxidation must mean what oxygen does to
- other atoms, which means it takes
- electrons away from them.
- Usually it takes 2 electrons for itself.
- Maybe it took 1 from each of the other atoms. So if you've
- been oxidized, you've had electrons taken away from you.
- And so you'll have a positive charge.
- Now, that interpretation is only partially true.
- The reality is it does not have to be oxygen that's doing
- the oxidation.
- So for example, let me do hydrogen fluoride.
- H F.
- I've shed my habit of writing fluorine as Fl.
- I've now remembered that its elemental symbol is just F.
- So hydrogen fluoride, if it was in an aqueous solution,
- it'd be hydrofluoric acid.
- That's hydrogen here bonding with fluorine, one of the most
- electronegative atoms. So what's going to happen here?
- Once again, the reality is that hydrogen is sharing--
- it's a covalent bond with fluorine.
- But the electron spends most of its time here, on the
- fluorine atom.
- So you're going to have a partial negative charge here,
- and a partial positive.
- But we don't like this partial, halfway game.
- We want to say, look.
- If this was, hypothetically, an ionic bond, if one of these
- people have to gain or lose an electron, how
- would it play out?
- In that situation, this guy likes to hog electrons more
- than hydrogen does.
- So the fluorine would gain an electron and have an oxidation
- number of minus 1.
- And the hydrogen would lose an electron, and have an
- oxidation number of plus 1.
- In this case, hydrogen has been oxidized.
- And notice, there's no oxygen to be seen.
- So the way I think about it, fluorine did to hydrogen what
- oxygen would have done.
- For example.
- If I say that you've been Bernie Madoff'ed.
- Bernie Madoff might not be the actual individual who's taking
- your money and putting it into a Ponzi scheme, but it would
- mean that someone else is doing the same thing that
- Bernie Madoff would have done to you.
- So in the same way, even though there's no oxygen here,
- fluorine has oxidized the hydrogen.
- Now, I'll introduce another word, and that's reduction.
- Reduction is the opposite of oxidation.
- I'll write it in blue.
- And this just means a reduction of charge.
- So in this hydrogen fluorine molecule, or hydrogen fluoride
- molecule, hydrogen has been oxidized.
- It's been oxidized because electrons have been
- taken away from it.
- That's what oxygen would have done to it.
- And so it has, now, a positive charge.
- And its oxidation number is plus 1.
- Because exactly one electron would have been taken away
- from this in a hypothetical ionic bond.
- Now, fluorine has been reduced.
- Its oxidation state was reduced by 1.
- If there was no hydrogen around, it
- would've been neutral.
- But now it has a minus 1 charge.
- Because in our hypothetical world of there's no partial
- charge, it's more electronegative.
- It took the atom from the hydrogen.
- So it has a minus 1 charge.
- Its charge has been reduced.
- Its oxidation state has been reduced.
- Your oxidation state is just that hypothetical charge.
- And that's how I think of it.
- Reduction means reduction in your hypothetical charge.
- Oxidation means, what would oxygen have done do you?
- Which means it would've taken away electrons, which would
- mean you have a positive charge.
- But if you look in some textbooks, or some teachers,
- they'll give you a mnemonic.
- Leo the lion says, GER.
- I'll say, says, in small, because it's really irrelevant
- to the mnemonic.
- And this is just a way of remembering that Leo means
- Losing Electrons is equal to Oxidation.
- And that Gaining Electrons is equal to Reduction.
- Now, to me, when you first said, oh wait, I'm gaining
- electrons, but I'm getting reduced.
- What's getting reduced?
- What's getting reduced is your charge, because
- electrons are negative.
- You're gaining something with a negative charge.
- So that's where the reduction comes from.
- Oxygen comes from the fact that oxygen will normally make
- you lose electrons.
- It will normally take electrons away from you, even
- though oxygen might not have anything
- to do with the reaction.
- Now, we can look at the periodic table, and we can
- guess, in most molecules, what an oxidation state of a given
- atom will be.
- All of these guys are alkali earth metals.
- And this is really a review.
- It's kind of taught as something new.
- But we know all of these guys.
- They love to give up electrons, because that makes
- them stable.
- Because they have this one electron in their outer shell.
- So they tend to have an oxidation state of plus 1,
- which means that they tend to give away an electron.
- For example, if I write sodium chloride.
- In this case, the oxidation state is truly reflective of
- their charge.
- Its charge is equal to its oxidation state.
- It gave away an electron.
- And that's true for all of the alkali metals right here.
- Remember, I don't include hydrogen there.
- Because hydrogen is a little bit of a special case.
- It could have been thrown here on the periodic table.
- Because it has one electron in its outermost shell, which is
- just its only shell.
- But it's also happy to get two, and have a configuration
- like helium.
- So you can almost view it as it's very close to
- completing its shell.
- So it sometimes has alkali metal-type properties.
- And sometimes it has halogen properties, where it wants to
- gain electrons.
- So hydrogen-- Let's say hydrogen is bonding with one
- of these guys, right?
- So let's say you had lithium hydride.
- So in this case, lithium, is right here.
- It loves to lose its electrons.
- So it'll lose its electron.
- So it'll have an oxidation state of plus 1.
- And it will lose its electron to hydrogen.
- Because hydrogen is more electronegative than lithium.
- Because hydrogen, you give it 1 electron then it has
- electron configuration like helium.
- So hydrogen will be minus 1.
- In this situation, hydrogen bonds with people roughly on
- the left-hand side of the periodic table.
- So lithium-- just to review our terminology.
- Lithium was oxidized by the hydrogen.
- Hydrogen was reduced by the lithium.
- Its charge went down.
- Their oxidation number?
- Plus 1 for the lithium, minus 1 for the hydrogen.
- What is the total sum of an oxidation state
- for a neutral molecule?
- In this case, you add up the charges and you get, oh, it's
- equal to 0.
- And that's a big takeaway.
- Then, in general, if you have a neutral molecule-- let me
- write that down.
- I'll just say neutral compound.
- Then the oxidation states add to 0.
- And if you have a non-neutral compound, if you have a plus 1
- charge, the oxidation states of all of the molecules in
- your compound are going to add up to your charge.
- Well, let me give another example.
- So in this case, I had hydrogen
- bonding with these guys.
- In that case, hydrogen is the one that took the electron.
- Let's say I had OH minus.
- So hydrogen, in this case, when it's dealing with
- something that's not on the left-hand side, when it's
- dealing with a super electronegative atom,
- hydrogen's oxidation state is plus 1.
- Right?
- Now, oxygen's typical oxidation state, when it bonds
- with almost everything else-- and I'll give a special case
- in a few moments-- is minus 2.
- Because it normally takes two electrons from other things.
- We saw that with the water.
- Oxygen's oxidation state, in most cases, is minus 2.
- So even though it hasn't gained two electrons here, we
- would write its oxidation number as
- minus 2 for the oxygen.
- And then you add up the two oxidation states,
- minus 2 plus 1.
- Well, actually I'll take that back.
- Oxygen has gained an extra electron because this is
- already an ion.
- So in this case, this OH, oxygen has stolen an electron
- from someone else.
- Let me draw that.
- And then oxygen normally only has 1, 2, 3, 4-- so oxygen has
- completed its valence shell.
- So essentially, it took an extra electron
- from some place else.
- So it definitely does have a minus 2 charge in this case.
- Oxygen has a minus 2 charge.
- Hydrogen has a plus 1 charge, because this electron was
- taken from it by the oxygen.
- And so the total charge of this molecule is minus 2 plus
- 1, which is minus 1.
- And that's why you have a minus 1 charge on OH.
- So you add up the oxidation states, and you're going to
- get the charge of the atom under question.
- Now, we already saw that most of these guys have an
- oxidation state of plus 1.
- These guys here, the alkaline earth metals, have an oxygen
- state of plus 2, because they like to give away two atoms.
- We already saw hydrogen.
- If they bond with people here, they take the electrons.
- So they get reduced.
- So it has an oxidation state of minus 1.
- But if hydrogen bonds with these guys on the right-hand
- side, it gives away the electrons.
- So it would have an oxidation state of plus 1.
- We saw that with hydrogen fluoride.
- It would be the case with hydrogen
- chloride, hydrogen bromide.
- You saw the case with water.
- All of these cases, hydrogen is the one that's giving away
- the electrons.
- So it gets oxidized.
- And these guys get reduced, or their the oxidizing agent.
- Now, what about these guys over here?
- What about our halogens?
- Well, they all like to take electrons.
- So their typical oxidation state is minus 1.
- Oxygen's typical oxidation state?
- We just saw it's minus 2.
- And I'll show you an example, which is really one of the few
- cases where oxygen doesn't have a
- minus 2 oxidation state.
- And then there's a lot of these other elements, in
- between, that can have multiple
- different oxidation states.
- And we'll see that in a lot of examples.
- But if you know that hydrogen is plus or minus 1, if
- hydrogen is bonding with these guys, it's a plus 1. it gives
- away electrons.
- If it bonds with these guys, it's a minus 1.
- it takes the electrons.
- If you know oxygen tends to take two electrons, so it has
- a minus 2 oxidation state.
- And the halogens take one.
- If you know that, you can normally figure out all of the
- other people in the compound.
- So what I'll do is I'll leave you there for now.
- I realize I'm running over time.
- In the next video, we're going to do some maybe more
- difficult examples.
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