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When you're studying chemistry,

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you'll often see reactions.

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In fact, you always see reactions.

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For example, if you have hydrogen gas, it's a diatomic

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molecule, because hydrogen bonds with itself in the

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gaseous state, plus iodine gas, I2.

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That's also in the gaseous state.

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It's very easy to say, oh, you know, if you put them

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together, they're going to react and form the product.

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If you have 2 moles of hydrogen, 2 moles of iodine,

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so it's going to form 2 moles of hydrogen iodide.

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That's all nice and neat, and it makes it seem like it's a

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very clean thing that happens without much fuss, but we know

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that that isn't the reality.

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And we also know that this doesn't happen just instantly.

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It's not like you can just take some hydrogen and put it

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with some iodine and it just magically turns

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into hydrogen iodide.

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There's some process going on that these gaseous state

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particles are bouncing around, and somehow they must bounce

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into each other and break bonds that they were in before

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and form new ones.

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And that's what we're going to study now.

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This whole study of how the reaction progresses and the

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rates of the reactions is called kinetics, which is a

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very fancy word, but you're probably familiar with it

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because we've talked a lot about kinetic energy.

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Kinetics, which is just the study of the rate of

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reactions, how fast do they happen and how do they happen.

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So let's just in our minds come up with an intuitive way

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that hydrogen and iodine can combine.

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So let's think about what hydrogen looks like.

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So if we get our Periodic Table out, hydrogen's got one

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valence electron, so if they have two hydrogen atoms, they

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can share them with each other.

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And then iodine has seven valence electrons so if they

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each share one, they'd get complete as well, so let's

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just review that right now.

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So hydrogen, this hydrogen, might have one.

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Well, it will have one electron out there.

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And then you could have another hydrogen that has

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another electron out there.

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And then if they form a bond, they share this.

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This hydrogen can pretend like he has this electron.

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This hydrogen can pretend like she has that electron, and

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then they're happy.

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They both feel like they've completed their 1s shell.

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Same thing on the iodine side where you have two iodines.

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They both have seven valence electrons.

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They're halogens.

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You know that already.

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Halogens are the Group 7 elements, so

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they have seven electrons.

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This guy's got one here.

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This guy's got one here.

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If this guy can pretend like he's got that

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electron, he's happy.

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He has eight valence electrons.

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If this guy can pretend like he's got that one, same thing.

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So there's a bond right here, and this is why hydrogen is a

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diatomic molecular gas, and this is why

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iodine is the same.

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Now, when they're in the gaseous state, you have a

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bunch of these things that are moving around, bumping into

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each other.

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I'll do it like this, so the hydrogen might look

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something like this.

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The hydrogen has these two atomic spheres

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that are bonded together.

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They have these electrons in between that are

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keeping them bonded.

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The iodine might look something like this.

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It's a much bigger molecule where it's bonded

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together like this.

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It's also sharing some electrons in a covalent bond

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and everything's probabilistic.

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So in order for these two molecules to turn into this,

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somehow these bonds have to be broken and new

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bonds have to be formed.

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And what has to happen is that these guys-- there's a ton of

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these guys.

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I could draw a bunch of them, or I could copy and paste.

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So there's a bunch of hydrogen molecules around, and some of

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these iodine gas molecules around.

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So what has to happen in order for us to get the hydrogen

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iodide is they have to collide, and they have to

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collide in exactly the right way.

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So let's say this guy-- actually, I can show it.

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Let's say he's moving.

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This is neat.

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I'm just dragging and dropping.

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But he's moving.

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He has to hit this hydrogen molecule just right.

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And maybe just right, if he just happens to hit it and

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bounce at it with enough energy, then all of a sudden,

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let's say we get to this point right here.

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These electrons are going to say, hey, you know, it's nice

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to be shared this way.

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We're in a stable configuration.

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We're filling this 1s shell, but look at this.

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There's this iodine that's close by and

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they really want me.

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They're much more electronegative than me to

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hydrogen, so maybe they're kind of attracted here.

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They don't know whether they want to be here between that

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hydrogen and this right here between that, and so they kind

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of enter this higher energy state.

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And similarly, these guys, they say, hey, wouldn't it be

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nicer-- I don't have to be here.

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I could kind of go back home to my home atom if this guy

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comes in here, because then we're going to have eight

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valence electrons, and the same thing's happening here.

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And this complex right here, right when the collision

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happens, this is actually a state.

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It's the high energy state, or the transition

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state of the reaction.

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And this is called an activated complex.

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You know, I just drew it kind of visually, but you could

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draw it like this.

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So hydrogen has a covalent bond with another hydrogen.

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And then here comes along some iodine that has a covalent

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bond with some other iodine.

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But all of a sudden, these guys like to bond as well, so

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they start forming-- so there's kind of a little bit

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of an attraction on that side, too.

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So this is another way of drawing

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the activation complex.

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But this is a high energy state, because in order for

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the electron, the way you can think of it, to kind of go

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from that bond to this bond, or this bond to that bond, or

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to go back, they have to enter into a higher energy state.

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A less stable energy state than they were before.

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But they do that if there's enough energy, because you can

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go from both of these things separate.

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Let me just draw them separate.

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You have both of them separate.

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You have the hydrogen separate plus the iodine separate.

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They go to this, which is a higher energy state, but if

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they can get to that higher energy state, if there's

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enough energy for the collision and they have enough

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kinetic energy where they hit in the right orientation, then

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from this activated complex or this higher energy state, it

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will then go to the lowest energy state.

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And the lowest energy state is the hydrogen iodide.

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Whoops!

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I want to draw iodide and then the hydrogen.

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This is actually a lower energy state than this.

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But in order to get here, you have to go through a higher

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energy state.

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And I could do that with an energy diagram.

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So if we say that the x-axis is the progression of the

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reaction, and actually, we don't know how fast it's

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progressing, but you could kind of view it as time in

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some dimension, and let's say this is the potential energy.

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I don't want to draw thicker lines.

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This is the potential energy right there.

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Let me make this line thicker as well.

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This is the potential energy.

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So initially, we are at this reality, and you can kind of

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view it as the combined potential energy.

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So essentially, we start off here, and this

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is the H2 plus I2.

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And a lower potential energy is when we're in the hydrogen

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iodide, so this is the lower potential energy down here.

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This is the 2HI, right?

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But to get here, we have to enter this higher activation

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energy, where the electrons have to get-- they have to

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have some energy to kind of be able to at least figure out

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what they want to do with their lives.

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And so you have to add energy to the system.

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You don't always have to add it, but if it doesn't happen

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spontaneously, you're going to have to add some energy to the

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system to get to this activated state, right?

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So this is when we're at this thing right here.

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We're there, so some energy has to be in the system.

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And this energy, the difference between the energy

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we were at when we were just hydrogen molecules and iodine

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molecules, and the energy we have to get to get this

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activated state-- this distance right here-- this is

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the activation energy.

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If we're able to somehow put enough energy in the system,

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then this thing will happen.

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They'll collide with enough energy and bonds will be

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broken and reformed.

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Activation energy.

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Sometimes it's written is Ea, energy of activation.

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And in the future, we'll maybe do reactions where we actually

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measure the activation energy, but the important thing is to

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conceptually understand that it's there.

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That things just don't spontaneously go

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from here to here.

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And I won't go deeply into catalysts right now, but

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you've probably heard of the word catalyst, or something

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being catalyzed.

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And that's some other agent, some other

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thing in the reaction.

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So right now, we have H2 plus I2

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yielding 2 hydrogen iodides.

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Now, you could have a catalyst, and I'll

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just say plus C.

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And I actually don't know what a good catalyst would be for

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this reaction.

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And how a catalyst operates is it can actually operate in

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many, many different ways, so that's why I don't want to do

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it in this video.

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But what a catalyst is is something that doesn't change.

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It doesn't get consumed in the reaction.

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The catalyst was there before the reaction.

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The catalyst was there after the reaction.

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But what it does is it makes the reaction happen either

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faster, or it lowers the amount of energy for the

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reaction to happen, which is kind of the same thing.

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So if you have a catalyst, then this activation energy

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will be lower.

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What it does is it might be some molecule that allows some

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other transition state that has less of a potential energy

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so that you require less heat or less concentration of the

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molecules for them to bump into each other in the right

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direction to get to that other state, so you

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require less energy.

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So given how we understand how these kinetics occur, or these

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molecules interact with each other, what do you think are

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the things that will drive whether a

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reaction happens or not?

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We already know that if we have a positive catalyst,

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there's something called a negative catalyst that will

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actually slow down a reaction.

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But if we have a positive catalyst, obviously, it lowers

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the activation energy so this makes a reaction faster.

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More molecules are going to bump into each other just

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right to be able to get over this hump because the hump

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will be lower when you have a catalyst.

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Also, if you increase the concentration, right?

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If you increase your concentration of molecules, if

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the concentration goes up, then you just have more stuff

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to bump into each other, right?

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There's just the likelihood.

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Everything is probablistic.

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When people write these reaction equations, it all

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seems nice and simple and very clear, and it happens.

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But no, in the real world, you just have things bumping into

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each other.

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And when we do biology videos, it will be fascinating to talk

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about, because every biological process is really

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just a chemical process, and it's really just the byproduct

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of all of these things bumping into each other.

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And you can imagine, the more concentration you have of the

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things that need to bump into each other, the more likely

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you're going to get just that perfect bump and that perfect

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amount of kinetic energy for the reaction to happen.

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Actually, I'll make a little other note here.

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This reaction, you might say, OK, I have some-- let's see,

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I'm at this energy.

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How do I ever get over this?

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How does this ever react?

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Well, remember, in a gas, the kinetic energies of all of the

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molecules, they're not uniform.

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Some molecules will have higher kinetic energy; some

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will have lower.

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Temperature just gives you the average.

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So there's always some probability that two maybe

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high kinetic energy molecules will bump into

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each other just perfectly.

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Surpass the kinetic-- so they have enough kinetic energy to

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get into the activation state, and then they can go to the

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lower state, which is the hydrogen iodide.

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At all temperatures this will occur, but obviously if you

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increase the temperature, that reaction is more likely.

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So that's the other one.

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So, temperature.

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Temperature is probably the single biggest thing that will

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make the reaction happen faster.

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So all of these things, you want higher temperature,

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higher reaction.

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And then if you just want to think about the molecules

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itself, if you have molecules where their original bonds are

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weak, they're more likely to be able to interact.

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And there's other things you could talk about: the

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molecular shape, how available certain atoms are to interact

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with other atoms. That really becomes significant when we

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start going into biology.

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And the last one, and you probably realize this, is just

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the surface area.

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If you increase the surface area-- so we were just doing

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gas-gas interactions, which almost by definition have

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pretty good surface area interactions.

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But if the surface area goes up, then the reaction also

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goes up, the reaction rate.

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And how do you think about that?

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Well, think about the reaction of-- you know, we've done this

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multiple times.

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Sodium chloride-- solid-- so solid salt, plus liquid water,

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leads to sodium-- well, we could think of it a lot of

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different ways, but we could think of it as sodium ion

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aqueous plus chloride anions-- this is a cation and anion--

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aqueous, so it gets dissolved.

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And how does that happen?

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If you have a big block of ice-- no, not ice, of salt.

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I'll do salt in grey.

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If you have a big block of salt in there, so there's a

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bunch of sodium and chloride atoms in it.

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And you have water all around it, the water is only going to

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be able to interact with the surface molecules and slowly

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dissolve away the salt, slowly make polar bonds.

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These are actually polar dipole bonds with the

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different sodium or chloride ions.

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But if you were to break this up into smaller cubes, if you

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were to break it up or really crush it into really small

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pieces, then all of a sudden the surface area that the

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water molecules can interact with, it can actually interact

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with more of the sodium chloride, so the reaction will

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happen faster.

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So surface area, if you increase the surface area of

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interaction, then you'll also increase the reaction rate.

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If you're trying to do it with two fluids, what you could do

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is you can kind of spray one fluid into the other, so you

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have little droplets, so you also

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increase the surface area.

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So anyway, this is kind of an introduction to the idea of

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kinetics, but hopefully, it gives you a sense that these

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reactions-- and I want you to really think about

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chemistry this way.

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Not think about it is as, oh, it's just some formula I have

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to remember, that these really are bumps and bruises between

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atoms. It's probabilistic and it's messy.

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And we really have to think about what will make it more

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likely that these things collide in just the perfect

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way for the reactions to happen.