Reaction rates
Introduction to Kinetics Kinetics, activation energy, activated complex and catalysts.
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- When you're studying chemistry,
- you'll often see reactions.
- In fact, you always see reactions.
- For example, if you have hydrogen gas, it's a diatomic
- molecule, because hydrogen bonds with itself in the
- gaseous state, plus iodine gas, I2.
- That's also in the gaseous state.
- It's very easy to say, oh, you know, if you put them
- together, they're going to react and form the product.
- If you have 2 moles of hydrogen, 2 moles of iodine,
- so it's going to form 2 moles of hydrogen iodide.
- That's all nice and neat, and it makes it seem like it's a
- very clean thing that happens without much fuss, but we know
- that that isn't the reality.
- And we also know that this doesn't happen just instantly.
- It's not like you can just take some hydrogen and put it
- with some iodine and it just magically turns
- into hydrogen iodide.
- There's some process going on that these gaseous state
- particles are bouncing around, and somehow they must bounce
- into each other and break bonds that they were in before
- and form new ones.
- And that's what we're going to study now.
- This whole study of how the reaction progresses and the
- rates of the reactions is called kinetics, which is a
- very fancy word, but you're probably familiar with it
- because we've talked a lot about kinetic energy.
- Kinetics, which is just the study of the rate of
- reactions, how fast do they happen and how do they happen.
- So let's just in our minds come up with an intuitive way
- that hydrogen and iodine can combine.
- So let's think about what hydrogen looks like.
- So if we get our Periodic Table out, hydrogen's got one
- valence electron, so if they have two hydrogen atoms, they
- can share them with each other.
- And then iodine has seven valence electrons so if they
- each share one, they'd get complete as well, so let's
- just review that right now.
- So hydrogen, this hydrogen, might have one.
- Well, it will have one electron out there.
- And then you could have another hydrogen that has
- another electron out there.
- And then if they form a bond, they share this.
- This hydrogen can pretend like he has this electron.
- This hydrogen can pretend like she has that electron, and
- then they're happy.
- They both feel like they've completed their 1s shell.
- Same thing on the iodine side where you have two iodines.
- They both have seven valence electrons.
- They're halogens.
- You know that already.
- Halogens are the Group 7 elements, so
- they have seven electrons.
- This guy's got one here.
- This guy's got one here.
- If this guy can pretend like he's got that
- electron, he's happy.
- He has eight valence electrons.
- If this guy can pretend like he's got that one, same thing.
- So there's a bond right here, and this is why hydrogen is a
- diatomic molecular gas, and this is why
- iodine is the same.
- Now, when they're in the gaseous state, you have a
- bunch of these things that are moving around, bumping into
- each other.
- I'll do it like this, so the hydrogen might look
- something like this.
- The hydrogen has these two atomic spheres
- that are bonded together.
- They have these electrons in between that are
- keeping them bonded.
- The iodine might look something like this.
- It's a much bigger molecule where it's bonded
- together like this.
- It's also sharing some electrons in a covalent bond
- and everything's probabilistic.
- So in order for these two molecules to turn into this,
- somehow these bonds have to be broken and new
- bonds have to be formed.
- And what has to happen is that these guys-- there's a ton of
- these guys.
- I could draw a bunch of them, or I could copy and paste.
- So there's a bunch of hydrogen molecules around, and some of
- these iodine gas molecules around.
- So what has to happen in order for us to get the hydrogen
- iodide is they have to collide, and they have to
- collide in exactly the right way.
- So let's say this guy-- actually, I can show it.
- Let's say he's moving.
- This is neat.
- I'm just dragging and dropping.
- But he's moving.
- He has to hit this hydrogen molecule just right.
- And maybe just right, if he just happens to hit it and
- bounce at it with enough energy, then all of a sudden,
- let's say we get to this point right here.
- These electrons are going to say, hey, you know, it's nice
- to be shared this way.
- We're in a stable configuration.
- We're filling this 1s shell, but look at this.
- There's this iodine that's close by and
- they really want me.
- They're much more electronegative than me to
- hydrogen, so maybe they're kind of attracted here.
- They don't know whether they want to be here between that
- hydrogen and this right here between that, and so they kind
- of enter this higher energy state.
- And similarly, these guys, they say, hey, wouldn't it be
- nicer-- I don't have to be here.
- I could kind of go back home to my home atom if this guy
- comes in here, because then we're going to have eight
- valence electrons, and the same thing's happening here.
- And this complex right here, right when the collision
- happens, this is actually a state.
- It's the high energy state, or the transition
- state of the reaction.
- And this is called an activated complex.
- You know, I just drew it kind of visually, but you could
- draw it like this.
- So hydrogen has a covalent bond with another hydrogen.
- And then here comes along some iodine that has a covalent
- bond with some other iodine.
- But all of a sudden, these guys like to bond as well, so
- they start forming-- so there's kind of a little bit
- of an attraction on that side, too.
- So this is another way of drawing
- the activation complex.
- But this is a high energy state, because in order for
- the electron, the way you can think of it, to kind of go
- from that bond to this bond, or this bond to that bond, or
- to go back, they have to enter into a higher energy state.
- A less stable energy state than they were before.
- But they do that if there's enough energy, because you can
- go from both of these things separate.
- Let me just draw them separate.
- You have both of them separate.
- You have the hydrogen separate plus the iodine separate.
- They go to this, which is a higher energy state, but if
- they can get to that higher energy state, if there's
- enough energy for the collision and they have enough
- kinetic energy where they hit in the right orientation, then
- from this activated complex or this higher energy state, it
- will then go to the lowest energy state.
- And the lowest energy state is the hydrogen iodide.
- Whoops!
- I want to draw iodide and then the hydrogen.
- This is actually a lower energy state than this.
- But in order to get here, you have to go through a higher
- energy state.
- And I could do that with an energy diagram.
- So if we say that the x-axis is the progression of the
- reaction, and actually, we don't know how fast it's
- progressing, but you could kind of view it as time in
- some dimension, and let's say this is the potential energy.
- I don't want to draw thicker lines.
- This is the potential energy right there.
- Let me make this line thicker as well.
- This is the potential energy.
- So initially, we are at this reality, and you can kind of
- view it as the combined potential energy.
- So essentially, we start off here, and this
- is the H2 plus I2.
- And a lower potential energy is when we're in the hydrogen
- iodide, so this is the lower potential energy down here.
- This is the 2HI, right?
- But to get here, we have to enter this higher activation
- energy, where the electrons have to get-- they have to
- have some energy to kind of be able to at least figure out
- what they want to do with their lives.
- And so you have to add energy to the system.
- You don't always have to add it, but if it doesn't happen
- spontaneously, you're going to have to add some energy to the
- system to get to this activated state, right?
- So this is when we're at this thing right here.
- We're there, so some energy has to be in the system.
- And this energy, the difference between the energy
- we were at when we were just hydrogen molecules and iodine
- molecules, and the energy we have to get to get this
- activated state-- this distance right here-- this is
- the activation energy.
- If we're able to somehow put enough energy in the system,
- then this thing will happen.
- They'll collide with enough energy and bonds will be
- broken and reformed.
- Activation energy.
- Sometimes it's written is Ea, energy of activation.
- And in the future, we'll maybe do reactions where we actually
- measure the activation energy, but the important thing is to
- conceptually understand that it's there.
- That things just don't spontaneously go
- from here to here.
- And I won't go deeply into catalysts right now, but
- you've probably heard of the word catalyst, or something
- being catalyzed.
- And that's some other agent, some other
- thing in the reaction.
- So right now, we have H2 plus I2
- yielding 2 hydrogen iodides.
- Now, you could have a catalyst, and I'll
- just say plus C.
- And I actually don't know what a good catalyst would be for
- this reaction.
- And how a catalyst operates is it can actually operate in
- many, many different ways, so that's why I don't want to do
- it in this video.
- But what a catalyst is is something that doesn't change.
- It doesn't get consumed in the reaction.
- The catalyst was there before the reaction.
- The catalyst was there after the reaction.
- But what it does is it makes the reaction happen either
- faster, or it lowers the amount of energy for the
- reaction to happen, which is kind of the same thing.
- So if you have a catalyst, then this activation energy
- will be lower.
- What it does is it might be some molecule that allows some
- other transition state that has less of a potential energy
- so that you require less heat or less concentration of the
- molecules for them to bump into each other in the right
- direction to get to that other state, so you
- require less energy.
- So given how we understand how these kinetics occur, or these
- molecules interact with each other, what do you think are
- the things that will drive whether a
- reaction happens or not?
- We already know that if we have a positive catalyst,
- there's something called a negative catalyst that will
- actually slow down a reaction.
- But if we have a positive catalyst, obviously, it lowers
- the activation energy so this makes a reaction faster.
- More molecules are going to bump into each other just
- right to be able to get over this hump because the hump
- will be lower when you have a catalyst.
- Also, if you increase the concentration, right?
- If you increase your concentration of molecules, if
- the concentration goes up, then you just have more stuff
- to bump into each other, right?
- There's just the likelihood.
- Everything is probablistic.
- When people write these reaction equations, it all
- seems nice and simple and very clear, and it happens.
- But no, in the real world, you just have things bumping into
- each other.
- And when we do biology videos, it will be fascinating to talk
- about, because every biological process is really
- just a chemical process, and it's really just the byproduct
- of all of these things bumping into each other.
- And you can imagine, the more concentration you have of the
- things that need to bump into each other, the more likely
- you're going to get just that perfect bump and that perfect
- amount of kinetic energy for the reaction to happen.
- Actually, I'll make a little other note here.
- This reaction, you might say, OK, I have some-- let's see,
- I'm at this energy.
- How do I ever get over this?
- How does this ever react?
- Well, remember, in a gas, the kinetic energies of all of the
- molecules, they're not uniform.
- Some molecules will have higher kinetic energy; some
- will have lower.
- Temperature just gives you the average.
- So there's always some probability that two maybe
- high kinetic energy molecules will bump into
- each other just perfectly.
- Surpass the kinetic-- so they have enough kinetic energy to
- get into the activation state, and then they can go to the
- lower state, which is the hydrogen iodide.
- At all temperatures this will occur, but obviously if you
- increase the temperature, that reaction is more likely.
- So that's the other one.
- So, temperature.
- Temperature is probably the single biggest thing that will
- make the reaction happen faster.
- So all of these things, you want higher temperature,
- higher reaction.
- And then if you just want to think about the molecules
- itself, if you have molecules where their original bonds are
- weak, they're more likely to be able to interact.
- And there's other things you could talk about: the
- molecular shape, how available certain atoms are to interact
- with other atoms. That really becomes significant when we
- start going into biology.
- And the last one, and you probably realize this, is just
- the surface area.
- If you increase the surface area-- so we were just doing
- gas-gas interactions, which almost by definition have
- pretty good surface area interactions.
- But if the surface area goes up, then the reaction also
- goes up, the reaction rate.
- And how do you think about that?
- Well, think about the reaction of-- you know, we've done this
- multiple times.
- Sodium chloride-- solid-- so solid salt, plus liquid water,
- leads to sodium-- well, we could think of it a lot of
- different ways, but we could think of it as sodium ion
- aqueous plus chloride anions-- this is a cation and anion--
- aqueous, so it gets dissolved.
- And how does that happen?
- If you have a big block of ice-- no, not ice, of salt.
- I'll do salt in grey.
- If you have a big block of salt in there, so there's a
- bunch of sodium and chloride atoms in it.
- And you have water all around it, the water is only going to
- be able to interact with the surface molecules and slowly
- dissolve away the salt, slowly make polar bonds.
- These are actually polar dipole bonds with the
- different sodium or chloride ions.
- But if you were to break this up into smaller cubes, if you
- were to break it up or really crush it into really small
- pieces, then all of a sudden the surface area that the
- water molecules can interact with, it can actually interact
- with more of the sodium chloride, so the reaction will
- happen faster.
- So surface area, if you increase the surface area of
- interaction, then you'll also increase the reaction rate.
- If you're trying to do it with two fluids, what you could do
- is you can kind of spray one fluid into the other, so you
- have little droplets, so you also
- increase the surface area.
- So anyway, this is kind of an introduction to the idea of
- kinetics, but hopefully, it gives you a sense that these
- reactions-- and I want you to really think about
- chemistry this way.
- Not think about it is as, oh, it's just some formula I have
- to remember, that these really are bumps and bruises between
- atoms. It's probabilistic and it's messy.
- And we really have to think about what will make it more
- likely that these things collide in just the perfect
- way for the reactions to happen.
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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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