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We now know what an oxidation state or an oxidation number

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is and what it means when things

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are oxidized or reduced.

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So let's see how that actually happens in reactions.

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So what we're going to study in this video is oxidation

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slash reduction reactions.

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All that is is a reaction where somebody's being

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oxidized, which means that electrons are being taken away

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from them, and someone's being reduced, which means that

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they're being handed electrons, or they're taking

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electrons away from someone else.

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Sometimes this has been termed-- because you have the

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R-E-D in reduction and you have the ox in oxidation, and

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they switch them around and they call this a redox

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reaction, which sounds like a very fancy chemistry term, but

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it just means a reaction where something is getting oxidized.

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And if something is getting oxidized, something else is

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getting reduced.

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So let's study a couple of them.

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So right here, I actually have combustion.

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This is actually methane.

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So methane is a fuel.

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You can make a motor powered on methane right there.

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It's a hydrocarbon.

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Most fuels that we use are hydrocarbons, which just mean

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carbons bonded in a bunch of different ways to a bunch of

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different hydrogens.

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If you add enough heat for the reaction to happen-- so its

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activation energy you have to put it in there with some

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oxygen around-- it's going to combust and it's going to

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produce carbon dioxide and water-- and I didn't put it

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here-- and even more heat than you put into it.

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So it's actually an exothermic reaction.

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It produces more heat than you put into it.

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I'll do a lot more of that in future videos on endothermic

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and exothermic reactions.

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But anyway, we care about the oxidation and the reduction.

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So let's see if anyone's getting oxidized, or anyone is

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getting reduced here.

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So let's look at their oxidation

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numbers, or their states.

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Here, carbon is bonded to 4 hydrogens.

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Who's giving, who's taking?

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So let's go to our Periodic Table.

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Carbon is here.

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Hydrogen is here.

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Carbon is more electronegative.

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It's adjacent to the Three Musketeers of

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electronegativity.

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These guys are the most electronegative.

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We always ignore the noble gases because they pretty much

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don't react at all.

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They're always happy with their 8 valence electrons.

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These guys love to get electrons.

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They're small molecules.

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Their outermost shell is close to the nucleus.

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They're so close to becoming noble that they just love

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hogging electrons.

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Carbon's almost there.

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Carbon's much further to the right on the periodic table

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than the hydrogen.

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So if you have carbon bonding with hydrogen, carbon is going

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to be the electron hog.

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So in the situation, let's go down here.

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So if carbon's the electron hog and hydrogen is having its

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electrons taken away from it-- remember, this is all kind of

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hypothetical.

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It's more partial, is the reality.

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But if you had to pick or choose, hydrogen's going to

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lose an electron each.

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So it's going to have an oxidation state of plus 1 for

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each hydrogen.

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You have 4 hydrogens each giving up an electron.

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So the carbon must be taking 4 electrons.

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So its oxidation number is minus 4.

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It's taken 4 electrons so its charge will go down by 4.

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So that's why it's negative.

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Fair enough.

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What's the oxidation state of this oxygen right there?

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Well, it's just bonded to itself.

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There's no reason to believe that 1 oxygen should be able

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to take any electrons from another oxygen.

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So it has a 0 oxidation state.

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It's not hogging more than its fair share of electrons that

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it was essentially born with.

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Now.

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After the combustion occurs, what are

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the oxidation numbers?

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Well, I have 2 oxygens bonded with carbon.

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Now we know oxidation was, to some

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degree, named after oxygen.

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It is one of the most electronegative.

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Almost anything bonded with oxygen is going to be giving

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up its electrons.

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We also know that oxygen likes to take 2 electrons.

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If it takes 1, it gets here.

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If 2, it gets here.

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It has a configuration like it has 8 valence electrons.

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So its typical oxidation number is minus 2.

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So carbon, in this in the situation, each oxygen is

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going to have a minus 2 oxidation number.

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The carbon dioxide molecule is neutral.

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So everything has to add up to 0.

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The oxygens, you have 2 of them, each of them with a

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minus 2 oxidation number.

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So that's minus 4.

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So carbon must be plus 4, which means that it has given

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up 4 electrons, right?

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Really it only has 4 electrons to give up.

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It has 1, 2, 3, 4 valence electrons in its second shell,

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which is its reactive shell.

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So its oxidation number is plus 4.

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Now let's look at the water.

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We've done about a bunch.

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When the hydrogens each give up an electron, they have an

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oxidation number of 1.

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Each oxygen takes 2 electrons.

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There's only 1 of them.

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So it's minus 2.

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So what's going on here?

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What's getting oxidized, what's getting reduced?

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You have the carbon.

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It went from an oxidation number of minus 4 an oxidation

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number of plus 4.

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So let me just draw what's happening to the carbon.

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Carbon goes from minus 4, which means it's hogging 4

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electrons, to a situation where it's having 4 electrons

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being hogged from it.

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It's kind of giving away 4 electrons.

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So someplace in this process, this guy must have given away

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8 electrons.

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This is a difference of 8 electrons.

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So plus 8 electrons.

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So what's happened to carbon?

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Carbon, has it been oxidized or reduced?

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Its charge has gone up.

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So it hasn't been reduced.

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Or the other way is that you can say electrons have been

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taken away from it.

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So it has been oxidized.

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Now, let's look at the oxygen itself.

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Oxygen, over here you have 4 oxygen molecules.

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So I'll just write 4 oxygen molecules here.

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And they all have an oxidation state of 0.

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From an oxidation-number point of view, they're neutral.

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On the side, we have 2, and then we

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have 4 oxygen molecules.

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What's their oxidation state?

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They're all minus 2.

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So each of these oxygens must have gained an electron.

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So actually, let me let me re-write this reaction.

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Let me erase a little bit of it.

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Even better, let me just move this over.

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There you go.

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Let me fill in that, so aesthetically it's pleasing.

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There you go.

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All right.

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So we have 4 oxygens with a 0 oxidation state turning into 4

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oxygens with a minus 2 oxidation state.

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So each of these 4 oxygens took 2 electrons.

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There's 4 of them, so we must have gained 8 electrons.

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So what happened to oxygen?

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It's oxidation state went down, or its hypothetical

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charge went down, or its charge was reduced.

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So this is reduced oxygen.

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What was it reduced by?

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It was reduced by the carbon.

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What was the carbon oxidized by?

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It was oxidized by the oxygen, which oxygen tends to do.

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It tends to oxidize things.

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What is the oxidizing agent?

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Well it's the thing that did the oxidizing.

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So the oxygen is the oxidizing agent.

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What is the reducing agent?

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Well it's the thing that did the reducing?

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It's the reducing agent, is the carbon.

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So you see in this redox reaction, carbon was oxidized.

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It gave away 8 electrons from this state to that state,

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hypothetically, and oxygen was reduced.

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Its charge was reduced by a total of 8 electrons, 2 for

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each oxygen.

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So its oxidation number went down.

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Let's do a couple more.

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Each of the things that I wrote-- just to kind of do a

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side-note-- this is called a half reaction, because I'm

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writing just what happened to the carbon.

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Here I'm writing just what happened to the oxygen.

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I did ignore something.

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I did ignore the hydrogen.

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I encourage you to do something like

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this with the hydrogen.

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But you'll see that hydrogen was

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neither oxidized or reduced.

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On both sides of the equation, all the hydrogens had a plus 1

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oxidation state.

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Let's do another one.

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Here I have another combustion situation.

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This is what happened to the Hindenburg.

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They filled a balloon full of hydrogen because they wanted

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it to float, because hydrogen is a very light gas.

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Unfortunately, there must have been a spark in the presence

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of oxygen, and it combusted.

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And actually, this is what is used for a

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rocket fuel as well.

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If you have liquid hydrogen-- actually, I don't think they

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have liquid hydrogen.

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Well, I don't know enough about rocket science.

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I'll have to do another video that in the future.

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But let's look at the oxidation states.

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What's the oxidation state of hydrogen here, in its

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elemental form?

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It's 0.

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2 hydrogens bonded to each other, no reason why they

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should hog or be hogged by another hydrogen.

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2 oxygen molecules or atoms bonded to each other?

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Once again, 0 oxidation state.

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Once they combust and form water, what's

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their oxidation states?

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We've done this multiple times.

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The oxygen has a minus 2 oxidation state.

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Each hydrogen has a plus 1.

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So let's write the half reactions.

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We have 2 hydrogens that are just happy as they are in

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neutral state.

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They end up being 2 H2's, or we could have written 4

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hydrogens, either way.

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They turn into 2 hydrogen molecules, because there's

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actually 4 hydrogen atoms, with a plus 1 oxidation state.

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So they must have each given away 1 electron.

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Now, how many total hydrogens are there?

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There are 4.

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So we must have given away 4 electrons.

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So this is a half reaction for hydrogen.

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Now.

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Let's do the same thing for our oxygen.

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We have some oxygen here.

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On the left-hand side, it has a neutral oxidation state.

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Then we end up with 2 oxygens on the right-hand side.

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I'll write it like this, 2O, each with the minus 2

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oxidation states.

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So each of these oxygens must have gotten 2 electrons.

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So it got 2 electrons each, or it gained 4 electrons.

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So what was oxidized?

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Oxidized means that electrons were taken away from you.

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The hydrogen was oxidized, oxidized by oxygen.

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What was reduced?

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The oxygen was reduced by the hydrogen.

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If you actually add these two reactions up, if you make this

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the left-hand side of your equation and you make this the

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right-hand side of your equation, you can say, OK,

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let's remove the electrons from both sides.

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Then you'll end up with your original reaction.

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Just to make sure our terminology is right, what's

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our oxidizing agent?

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It's the thing doing the oxidizing.

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It's the oxygen.

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What's our reducing agent?

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It's the hydrogen.

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Let's do one more.

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So here I have iron plus some hydrochloric acid.

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Let's say this is in an aqueous solution.

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You end up with iron (II) chloride plus some hydrogen.

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So let's do some oxidation numbers.

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I'll do it fast this time.

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Iron is just by itself.

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It's got a 0 oxidation state.

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Hydrogen with chloride, chloride's a halogen.

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These guys love to take electrons.

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They love to take 1 electron.

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They typically have a minus 1 oxidation state.

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So the chlorine is going to have minus 1.

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The hydrogen is going to be plus 1.

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You add them together, you get to a neutral compound.

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Fair enough.

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Now you go on this side.

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What is chlorine's oxidation state?

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Well, once again, it always likes to take 1 electron.

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So this is minus 1.

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But I have 2 chlorines here.

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This is a neutral compound.

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So the iron oxidation state must be plus 2.

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What about the hydrogen?

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Now it's just in its elemental form so

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it's going to be neutral.

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So it has a 0 oxidation state.

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So what was oxidized?

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Our iron went from neutral to having 2 electrons taken away

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from it, which may give it a positive charge.

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So plus 2 electrons got taken away.

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So this was oxidized.

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What about the hydrogen?

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The hydrogen went from 2 hydrogens with a plus 1 state,

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and essentially, they went to 2 hydrogen

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with a neutral state.

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So 2 electrons must have been added to the hydrogens.

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Their charge was reduced.

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So they were reduced.

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What were they reduced by?

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They were reduced by the iron.

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What was the iron oxidized by?

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It was oxidized by the hydrogen.

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What about the chlorine?

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The chlorine has a minus 1 oxidation number here.

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It has a minus 1 oxidation number here.

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It was a neutral relative to the redox reaction.

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Anyway, I think you get the point now.

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You can do this with a bunch of reactions.

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But this will give you a little bit more insight of

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actually what's going on, and who's

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gaining or losing electrons.

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In the future, it will also help us think about a lot of

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the structures of molecules.

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Anyway, see you in the next video.