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1.08 Collision Theory

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Advanced Chemistry

University of Kentucky

4.7 (419 ratings) | 50K Students Enrolled

A chemistry course to cover selected topics covered in advanced high school chemistry courses, correlating to the standard topics as established by the American Chemical Society. Prerequisites: Students should have a background in basic chemistry including nomenclature, reactions, stoichiometry, molarity and thermochemistry.

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KO

Jul 17, 2017

Excellent content and delivery!\n\nThis course has many well-explained examples and practice questions to reinforce the understanding of the taught concepts. Thank you!

CW

Sep 17, 2015

I used this course as a significant portion of my preparation for the CLEP Chemistry exam, and passed today with an A! Thank you!

From the lesson

Kinetics

The study of chemical kinetics is the study of change over time. It answers questions like: How fast are reactants consumed? How fast are products formed? This unit is dedicated to the exploration of how these questions are answered. We will look at the experimental evidence of how concentration affects these rates. We will also examine what occurs on the molecular level, especially with respect to the motion of molecules, that affects rates of reactions.

1.01 The Rate of Chemical Reactions10:12

1.02 Comparing Rate of Change for Reactants and Products13:52

1.02a Obtaining a Rate Law from Experimental Data Equations4:23

1.03 The Rate Law9:01

1.04 Obtaining a Rate Law from Experimental Data16:38

1.04a Rate Law Calculations8:20

1.05 First-Order Kinetics and the Integrated Rate Law0:00

1.05a Graphic 1st Order7:10

1.06 First-Order Kinetics and the Half-Life9:32

1.07 Second-Order Reactions8:18

1.07a Graphics 2nd Order6:02

1.08 Collision Theory13:13

1.08a Activation Energy3:48

1.09 The Arrhenius Equation15:32

1.09a Graphic Activation Energy2:57

1.10 Reaction Mechanisms12:45

1.10a RDS Fast Equilibrium4:38

1.10b Rate Determining Step2:41

1.11 Catalysis4:23

Taught By

Dr. Allison Soult
Lecturer
Dr. Kim Woodrum
Sr. Lecturer

Transcript

Up to this point we've been looking at macroscopic properties and

observations of reactions.

Kind of from the outside looking in, looking at concentration,

time rates of reactions.

We're gonna look at one other macroscopic proprerty of temperature and

see how temperature affects the rate of reaction here in just a moment.

But our learning objective for

this lesson is to be able to use something called Collision Theory.

So that's what we're talking about.

And Collision Theory is looking at things, what's occurring on the molecular level?

Down there with those molecules as they collide,

therefore the name Collision Theory.

And we're gonna use Collision Theory to recognize the reason for

the dependence of the rate of the reaction on concentration, which we've looked at in

the past, as well as on temperature which we're gonna observe here in just a moment.

I like to bake and most cake recipes occur at 350 degrees fahrenheit oven.

That's a good hot oven, but not extremely hot.

But there are some recipes that cook at a cooler oven.

And I've got this pound cake recipe that I bake the cake at 250 degrees.

Now in terms of what's the difference, what's the difference in time?

My pound cake recipe requires over an hour to cook,

whereas most cake recipes will take around 30 minutes to cook.

So a cake at 350 will cook faster.

And this is true for almost every reaction,

that the rate of the reaction is going to increase with temperature.

There's a few exceptions, and

we're not gonna get into why the exceptions occur at this moment.

But it's usually dealt with some kind of side thing that's going on,

alternate path that is taking place for the reaction.

If you think about things on the opposite end of the spectrum.

Let's say you've eaten already and you've got leftovers.

What do you wanna do, you wanna slow down a certain reaction,

that would be the rotting of the food, the growing of bacteria.

So you're always told not not leave that food out on the counter.

To put them away rapidly into the refrigerator, so

you slow down that process.

You want to keep it even longer?

You put it in the freezer.

And you can keep it for weeks and months.

If you've got a deep freeze, that means you're freezing it at a really low,

low temperature, you can keep it for even longer.

So that is a macroscopic property.

We see that reaction rates decrease as temperature goes down,

increases as temperature goes up.

We see that reaction rates decrease as concentration goes down,

increases as concentration goes up.

And in this lesson we're gonna see what's happening on

the molecular level that causes those things to happen.

I wanna write a rate law, just a generic rate law here for reaction.

Rate is equal to k times concentration raised to some power.

We have seen that as concentration goes up, rate goes up.

And we've talked about this rate constant being a constant.

But only being a constant at a constant temperature.

Now if the rate is going to go up with temperature, where does that play out?

It has to play out in the rate constant k.

And we see that if we plot rate constant and

temperature, as the temperature goes up, the rate constant goes up.

It's not a linear relationship but

it is definitely an increase in rate constant with an increase of temperature.

Oh, went the wrong direction, bear with me.

Okay, so we're gonna start by thinking about what's happening on the molecular

level that causes the rate to increase with concentration.

And then we'll look at why it increases with temperature.

The first thing that we need to know about Collision Theory is that for

a reaction to take place, a molecule has to collide.

They collide with each other.

If molecules just sitting there minding it's own business,

it's not gonna want to turn into a new product.

So if you can get the number of collisions to increase,

you're gonna make the reaction rate increase.

It's gonna happen at a more rapid rate.

So rate is directly proportional to the number of collisions per second.

So the first thing we can think about is,

why does rate go up as the concentration goes up?

Okay, what happens to the number of collisions as the concentration increases?

Well let's look at this little schematic of reaction.

The blue has to run into the red for a reaction to occur, let's say.

And the number of the spheres you see is a representation of the concentration.

Now if blue has to collide with red,

the black lines represent the number of possible collisions that can occur,

and we see four lines here, four possibilities of collisions.

Now if in the same size container, we increase a number of the blue spheres,

then we can increase the number of possible collisions to six.

So we're going to have a greater probability of these collisions occurring

in a certain amount of time, if we've just increased the blue.

We could increase the red, and this again would be six,

we've increased the red instead of the blue.

But if we increase both of them, then this is going up to nine here.

We have nine possible collision paths that would lead to a reaction.

So as you increase the concentration,

you're gonna increase the number of possible collisions.

So that we know rate increases with concentration this is why.

Rate is increasing with concentration because you are increasing

the number of collisions.

Now let's think about it in terms of something that you can

imagine a little bit better maybe.

If I put you in a room with three other people,

there's four of you all together running around the room.

And I say okay, I want you all to run in this room.

And if you run into something, you just bounce off and keep on running.

And you just imagine in your mind how many times you probably would run into somebody

as you're running randomly around the room.

Let's keep the room size the same, but

let's put a hundred of you running around in the room.

What is it going to do to the number of collisions?

If the number of you in the size of a room that would be a representation of

molarity, moles per liter.

As you increase the concentration,

you're certainly increasing the number of collisions.

So if you have to have a collision for a reaction to occur,

if you increase a concentration you're gonna increase the number for collision.

Therefore you're going to increase the rate of the reaction.

Okay, so that's the first thing, you have to have a collision.

But not every collision results in a reaction.

If it did, then all of our reactions would occur virtually, instantaneously.

In the blink of an eye, in a snap of your fingers the reaction would be done

because molecules move very, very rapidly, way faster then you can run around a room.

And there's so many collisions in such a short amount of time,

that the reaction would take place, virtually, instantaneously.

So why do some reactions take minutes, and some hours, and

some even maybe years to take place?

Well the reason for this is something that we need to understand.

And that is a collision has to occur with

enough kinetic energy to break the bonds that are there.

Those atoms are held together with energy.

We have to break those bonds and that requires energy to do.

If they don't collide with enough kinetic energy to actually break the bonds that

are there already, then they just bounce off of each other.

Activation Energy is the minimum energy required

to initiate that chemical reaction.

It is that barrier.

Anything less than that, molecules bounce off.

Anything there or more, the reaction can take place.

These diagrams are called Potential Energy Diagrams.

Let me just explain a few things about them first.

The one on the left is an exothermic reaction and

the one on the right is an endothermic reaction.

In your past lessons, somewhere in your past, you've talked about endothermic and

exothermic and you've talked about delta H.

The difference in the energy between products and

reactants would be the measure of this delta H.

For exothermic reactions, it would be products energy which is a small

number the reactants energy which would be a large number.

And these values are negative.

But as time goes by, as a reaction progresses from left to right.

You start with reactions, you're gonna go to products.

If there is enough energy to break those existing bonds, and

you have met that activation energy.

So I will call this the mountain those reactants have to climb,

the energy that they have to have in order to get over there and down to products.

Now over on the right hand side this is an Endothermic reaction.

The delta H again would be from across my arrow here.

And this would be a large number minus a small number, and they're always positive.

But the activation energy barrier that has to be overcome always starts from so

progressing to the right here.

The arrow goes from reactants all the way to the top of the peak, there.

So this reaction has got a much higher activation energy barrier.

Molecules have to collide with a lot more energy to become products.

Up here at the top of the mountain is something called an activated complex.

And it's gonna be a temporary species.

It is there for such short time.

It's usually not even observable.

But it is the temporary species formed

by the reactant molecules as a result of that collision.

So you can think of it as the jumbled mess

of atoms as they have collided with enough energy to break those bonds.

Once that activated complex is formed, that molecule or

those atoms can either recombine and go over here and reform reactants.

Or recombine and go over here and recombine and form products.

And truth of the matter is, both of them are gonna occur, but

if the reaction gonna proceed to products we're gonna form C and D.

Okay, so let's go back to thinking about collision theory.

Molecules have to collide for the reaction to occur.

You get more molecules you get more collisions but

they also have to collide with enough energy to get over this mountain.

So it's gotta get over this mountain.

So if they collide and they don't have enough energy, they just are A and

B still.

But if they collide with enough to get over the mountain, then you form products.

So based on what we know, at this point,

why does the reaction rate increase with temperature?

Well I've got two diagrams here, and

temperature one is less than temperature two.

So the right hand diagram is a higher temperature.

And this is a Maxwell–Boltzmann distribution

of the number of molecules that have a certain speed.

And let's say that kinetic energy, The average

kinetic energy is equal to one half mass times this speed squared.

Average speed, okay?

We have to have enough kinetic energy to get over the mountain.

And let's say that every molecule that has at least this speed, or higher,

will have enough kinetic energy to overcome the activation energy barrier.

So everyone of these molecules, if they were to collide, could react and

form products, but we raise the temperature.

When we raise the temperature, this distribution changes and

this is still in minimum speed needed to obtain this kinetic energy so

that it can get over the mountain.

But now all of these molecules have enough energy to get over the mountain.

So if temperature increases,

you might increase the number of collisions a little bit.

But really what you're doing is you're giving those molecules more kinetic energy

and more of them have enough energy to get over the activation energy barrier.

So this is our look at collision theory and trying to explain what's happening

on the molecular level that causes the rate to increase with concentration.

Increase the concentration you have more collisions, and

causes the right to increase with temperature.

If you got a higher temperature,

you have more molecules that can get over that mountain.

That activation energy barrier to get to the other side and form products.

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