>> Welcome to Chemical Reactions, Part Five.
In this video, we are going to look at chemical equilibrium
and the Le Chatlier's Principle.
And this is very important because as a biological system,
we are in constant equilibrium and our equilibrium changes
and this so will be an essential topic
for understanding our health and how chemistry plays a roll.
Let's get started.
All right, so.
Equilibrium is a term when the reverse reactions proceed
at the same rate, right, then the concentration or reactants
and products will no longer change
and the reaction has reached a state of equilibrium.
Okay. So before this, before this chapter,
we were just looking at reaction arrows where we said it went
from the reactants to the, oops, products, and we would just say
as all one hundred percent complete.
Okay, and many reactions are like that.
But, these reactions are both the forward
and the reverse reaction can occur.
Right. So, then, we can have a mix of reactants
and products simultaneously
because our forward reaction is occurring at the same rate
as the reverse reaction when we're adding equilibrium.
Okay. So, let's focus on equilibrium first
and then we will look at what happens
when we disrupt equilibrium.
Okay. So, we will start with a really simple reaction.
We'll just take this compound HA and it's going to dissociate
into H plus and A minus.
This is actually a weak acid and we'll study that in the future.
So don't worry about the acid part right now,
just focus on the equilibrium.
All right.
So what happens, right, so initially,
let's say we have one hundred percent of HA.
Okay. And so, we have a, and I'll make a graph here.
All right.
So for the forward reaction, we'll use a green arrow
and for the reverse reaction, we'll use a red arrow.
Now, we're going to look right here, right.
So, this will be our reaction rate on the "Y" axis,
and this will be time on the "X" axis and so, what happens,
initially, all we have is HA, so it dissociates and so,
it has a very high rate, but then,
as the concentration decreases, the rate lowers,
and so this right here, right, this is going to be
when we're at equilibrium.
And the reason why, is because, initially,
there is no reverse reaction because there are no products.
But once HA starts to dissociate, then each plus
and minus will be formed and so then we'll start
to get our reverse reaction and there we go.
And so, right here, at this time, we see that the forward
and reverse are at the same reaction rate
and so then this would be,
right where the reaction reached equilibrium.
Okay, so it didn't start an equilibrium.
So, we just started with pure HA.
It started to dissociate and then as the H plus
and A minus formed, they wanted to reconnect
and so we reached equilibrium.
Okay. So, now we will try to line
up this next graph the same.
Okay. So that's our reaction rates.
So the reaction rates are the same at equilibrium.
Now, let's look at the reactant and product concentration.
Okay, so, we'll keep HA as green.
Here we go with our forward reaction, okay.
And I am want to, I'm going to try to line this
up so we have it all happening together.
So, we'll go here.
All righty.
So now, along our "Y" axis, we have the concentration, right,
of the reactant, or product, whatever we're looking at.
So, here we're going to look at the reactant.
And so, what we will see is that, since we started
with pure, oops, I'm sorry, I have lost the page.
Since we started with pure HA, the concentration was high
and then it decreases and then
when we reach equilibrium, it remains unchanged.
And then at the same time, we have the A plus and the A minus,
right, so initially, there wasn't very much,
but then as the reaction proceeds, there we go.
So then, this flattens out.
So here we are.
We are at equilibrium.
At equilibrium.
So, at equilibrium, the rates of the forward
and reverse reactions are the same and the concentration
of reactants remains constant and the concentration
of the products is no longer changing.
So this should be, this red line, should be nice and flat.
Okay. So, if you look at this, all right,
we have to understand an equilibrium.
Okay. So, is the concentration, let's make sure
that this was our products, all right,
and this was our reactants.
So, this is a common misconception,
is the concentration of products equal to the concentration
of the reactants at equilibrium?
No. Right.
So, what we can say is that the concentration of reactants
and the concentrations of products remain constant.
Ooops! At equilibrium.
Okay. All right.
The forward and reverse rates are equal but the concentrations
of products and reactants can be different;
they just don't change.
So, at equilibrium, have the reactions stopped?
No. Right.
Because these concentrations are remaining the same,
it gives the illusion that the reactions have stopped.
But that's not the case.
Okay. So now, let's dig a little bit deeper into equilibrium.
Okay. So, it's often called Dynamic Equilibrium, right,
because of the reaction rates.
Right. It's a Dynamic Equilibrium
because the forward reaction, right, the forward
and reverse reactions, continue to occur.
Okay. So, that the unchanging concentrations of products
and reactants is an illusion
that there's nothing happening, right.
So, on the surface, it looks like nothing's happening,
but inside the forward and reverse reactions are happening.
Okay. So now, let us integrate everything that we've learned
about reactions to date.
All right.
So, we're going to do our reaction energy diagram.
So I'll give us some space for that.
So for a reaction energy diagrams, we are going
to have energy on our "Y" axis and we'll put time
on our "X" axis, okay.
And, so what type of curve?
So, we're going to do exothermic, right,
they're considered spontaneous
because the products are more stable than the reactants,
right, but spontaneous does not necessarily mean fast.
All right.
So, let's do this, so we'll put exothermic,
so our products are lower, more stable lower energy
than our reactants, and, of course,
if I gave you an actual reaction,
you would show the chemical formulas, okay.
And exothermic means, right, that this energy is going
to be released, so we put minus delta H, or delta H less
than zero, two ways to show that, okay.
Then, right here, right here is our activation energy,
our E sub a, so your reactants need to get up to,
what's this called up here?
Right? This is the transition state.
So, remember from collision theory, right,
that the reactants have to collide
with a minimum collision energy, the E sub a,
the activation energy, to, for a reaction to occur.
And not only do they have
to have this minimum collision energy, but they need
to have the correct orientation.
Okay. So, then, right, so this is exothermic,
it's considered spontaneous, but it's not necessarily fast
because it has a large activation energy.
And so then we have learned
that we can add a substance called a catalyst.
And what the catalyst does, right, so now with the catalyst,
we can lower the transition state energy, right,
which lowers the activation energy,
which gives us the faster reaction rate.
Okay. So, now let's move forward.
Have a little review tying this all together.
And so the main thing here with equilibrium, right,
especially if we have a low activation energy is we can have
the forward and the reverse reaction
happening simultaneously.
All right.
So now, let's go, we understand a system at equilibrium.
Now comes the more interesting thing.
Changing equilibrium because our body wants us to be
in equilibrium and then we're doing things
to throw it out of equilibrium.
We're exercising.
We're eating.
We're drinking water.
We're doing all kinds of things.
All right.
So, the gentleman by the name of Le Chatlier,
he came up with this principle that if we start with a system
at equilibrium and we add stress, we apply stress
to the system, the equilibrium is going to shift
to relieve the stress.
Okay. So, keep in mind, right, that reactants are on the left
and products are on the right.
Okay. So, we're going to talk about things going to the left
or the right and just remember that that's referring
to the reactants and the products.
Okay. So, you know, humans, we know a lot
about what human stress is like.
What causes stress in a chemical equilibrium situation?
Well, if we change the amount, right, the concentration amount,
or concentration, we use the square brackets
for concentration, right, of reactant, right, or product.
That's going to create stress, right.
If we change the temperature, that will create stress, or,
we can change the pressure or volume of a system,
and that will create stress.
Now with pressure and volume, the important thing to know,
this is for gas phase materials, reactants
and products only, gas phase only.
So we only worry about the effects of pressure and volume
when we have gas phase reactants or products.
Okay. So, when we're doing the stress, we can think about it,
so, here's our system at equilibrium.
Right. So, we can increase, right, or decrease things.
Now, what does that do to our, our balance?
It shifts our balance to the right, right?
So, if we cause stress that raises, right,
increases something on the reactant side or decreases
on the products side,
we say "we're shifting the equilibrium to the right".
Okay. We can also shift to the left, but I'm just wanted
to give you a little one-up before we get started.
All right.
So now we're going to look at an actual chemical reaction.
So, here we are at equilibrium.
So the phosphorus trichloride and chlorine are reacting
to produce the phosphorous pentachloride and some heat.
So this is the equilibrium and we start it
at equilibrium, right.
So add equilibrium to forward reaction, rate is equal
to the backward reaction rate and all
of the concentrations are remaining constant.
But notice they are not equal to each other, right.
Okay, so now, what we're going to do, now we add the stress.
Okay. So here is the stress.
We're going to add the chlorine, all right.
So here's the stress.
Okay. So, what did we do?
We increased the chlorine concentration.
So here is our teeter totter, right, the chlorine is over here
on the reactant side, right.
And what did we do?
We increased it.
We made it go up, right?
Increasing is going up.
So what have we done to our teeter totter or to our balance?
Right. We've shifted it right?
So, if we add the chlorine, we are going
to shift the equilibrium, right, to the right,
right, to the products side.
Okay. So, now the system has reached new equilibrium.
All right.
So, in theory, right, if we put the stress, we added stress
on the left, it pushed the reaction to the right,
so we would expect to have,
now we already expect to have more product.
Let's look and see, and we would expect to have less reactants.
So let's check this out.
Okay. So, at the new equilibrium, we compare,
so we add the original equilibrium,
we added the stress, and now equilibrium has been restored.
Okay. So at the new equilibrium, what happened
to the phosphorous trichloride?
Right? The concentration
of the phosphorous trichloride, right, it decreased.
What happened to the chlorine?
Right. Well, it went from ten to eight.
The concentration of the chlorine decreased two.
And then, our product, right, went from five to six.
So the concentration of our product increased, okay.
So when we say that we've shifted the equilibrium
to the right, it means that the reactants are going to react
with each other to produce more products.
So, which way did the equilibrium shift?
To the right.
Okay. So now let's try a different one.
Okay. So we start with our reaction and equilibrium
and now we are going to decrease the phosphorous trichloride.
Okay. So, the phosphorous trichloride,
it's on the reactant side.
Okay. So we always look to our reaction to figure out what side
to put the substance in question.
And now, we're going to decrease it.
Right. So, decreasing, we're going to go down.
So if we decrease, then that's going to cause our balance,
right, our teeter totter, to shift to the left.
All righty.
So, we would say that then we have shifted the equilibrium
to the left, right, which is the reactant side.
Okay. All righty.
Let's keep going because we talked
about there being more ways to stress a reaction,
so we've looked at changing concentrations of reactant
and our product, now we're going to look on the next pages
at what happens when we change temperature, pressure or volume.
So, let' bring in the next page.
So we'll just keep with the same reaction so we don't have
to worry about any of that.
All right.
So with temperature, right, this is where we're going
to take advantage
of understanding whether the reaction is endothermic
or exothermic, okay.
So, here heat is a product.
So we would describe this reaction, right, as exothermic,
all right, so the value of delta H, positive or negative, right,
so that means that delta H is negative,
which means it's less than zero.
Okay. All righty.
So now, we look at our reaction in equilibrium, right.
Heat is on the right, so we're going to put heat over here
on the right and so now, if we increase the temperature, right,
that means we're going to go up on the side so if we push
up on the right side, right,
we're going to shift our equilibrium,
right, shift to the left.
Okay. So, we can see that if we increase the temperature,
it favors the reverse reaction.
And, I've already answered this one;
if we increase the temperature, right,
the equilibrium shifts to the left.
All right.
Okay. Right.
So, if we increase the temperature,
what are we doing here?
I've missed this part.
Right. So, if we increase the temperature,
we are definitely adding heat.
Okay, so that's why we did that.
Okay. So, then, if we decrease the temperature, right, so heat,
this is an exothermic reaction.
Heat is on the product side.
Now, if we decrease the temperature, right,
we're going to pull down here, so we're going
to shift the equilibrium right [writing] Okay.
So, decreasing the temperature is actually going
to favor the forward reaction.
All right.
So, if we decrease the temperature,
we shift to the right and this is going to favor products.
Then, other part of the videos is just kind of making sure
that we can talk about this in a variety of terms
and we can ask questions in a variety of ways and so we want
to make sure, we don't get thrown off
by have a question as asked.
Okay, so, if we are going to do a reaction energy diagram
for this reaction, right, so now we are going to put
in a little more detail than in our previous one
because now we have an actual chemical species to talk about.
Right, so it's still exothermic.
So, the phosphorous pentachloride is a lower energy
than the phosphorous trichloride and the chlorine, right,
and so here's that heat released, right,
the delta minus H, there's that heat released
because it's exothermic, and then, of course, we want to put
in our activation energy and the transition state.
Okay, so, we're bringing it all together, so we've looked
at the effects of concentration.
We've looked at the effects of temperature.
And now, we are going to look more closely at the change
of pressure and volume.
Okay. So, here we go.
Voila. All right.
So, now, when we are changing pressure,
remember that we only look at pressure
and volume, only look at gases.
All right.
So, when we look here, we see that we have one more
of phosphorous trichloride reacting
with one mole of chlorine.
They're both in the gas phase.
So we see that we have two moles of gas on the reactant side
and when we look on the product side,
we only see one mole of gas.
Okay. So, I notice this pattern, so this is a benetism.
I haven't seen it anywhere else, but it's pretty helpful.
Basically, what happens then is,
let me make sure this is right, yeah, okay.
So, basically, we need to balance
out the volume, all right.
So, if we have two moles of gas here and we only have one mole
of gas here, we can put a V here for volume.
Okay. Now, from gas laws, we look at Boyle's Law, right.
We know that P one V one equals P two V two.
So, what that tells us is if we increase the pressure,
we're decreasing the volume, or vice versa.
If we decrease the pressure, we're increasing the volume.
So, there's a relationship between pressure and volume
that we have to bring to applying Le Chatlier's Principle
to gas phase reactions where we change pressure and volume.
So, that's why we talk about these together
because they are linked by Boyle's Law.
Okay. So, if we look here, right, so how many moles
of gas are present as reacting, we've already answered that?
Two moles.
How many moles of gas are present as product?
Only one mole.
Okay. So, if we look here, all right, so,
if we're increasing the pressure, right, then that means
that we're decreasing the volume.
So, notice that the V is on the right.
All right.
So, if we increase the pressure, then we're decreasing the volume
and so that is driving a reaction to the right.
Right. So that shifts the equilibrium to the right.
Okay. And so, let's think about this physically, right.
So, in the reactant side, we have two moles of gas clinging
for space and on the right side,
we only have one mole of gas wanting space.
All right, so as we increase the pressure, we're squeezing,
right, so we're squeezing the space and so, the reaction wants
to have, uhm, less moles of gas.
All right, so, as we increase the pressure,
we decrease the volume and that creates stress for the side
with more moles of gas.
Okay. And so from Le Chatlier's system,
the reaction wants to relieve stress.
So, it goes to where, aah, I don't need as much space.
That smaller volume is okay.
We don't need as much space.
Okay, let's look at another reaction.
What would happen if we change the pressure of this reaction?
Would there be any effect?
Let's look.
Right. So, we have one mole and one mole,
so we would have two moles of gas on the reactant side
and with our co-efficient here, we have two moles
of gas on the product side.
So, if we change the pressure or the volume, right, right, so,
changing pressure or changing volume, right, delta for change,
changing pressure or volume, has not effect, right,
because the same number of moles of gas are on each side.
Okay, so the other thing to keep in mind is
that these area all potential sources of stress
within a reaction, but, uhm, they don't necessarily, look,
are relevant to every reaction you study, okay.
So, let's look at our last page here and tie in one last thing.
We've talked about catalysts.
So, taking into consideration everything that we know
about reactions so far, is a catalyst going
to change the position of the equilibrium?
No. Right?
What does a catalyst change?
A catalyst change, right,
the catalyst decreases the activation energy
which increases the reaction rate.
So, basically, what a catalyst does, is it,
it lets the reaction get to equilibrium faster,
but it doesn't change the position of equilibrium.
Right, so it allows equilibrium
to be reached more quickly, faster.
Okay. But, it doesn't change the destination.
It's like putting in the express way.
Right. And so, we only just talked about this, right.
So, here would be our normal reaction.
Here, energy on the right, time on the X, right.
Here is without catalyst, okay.
And then if we add the catalyst with catalyst, okay.
So here, there's our delta minus H. Okay, look,
this energy difference between reactants and products, right,
this determines the equilibrium.
Okay. So the relative energy of the reactants and products,
oops, is going to determine our equilibrium and remember,
that the catalyst
or the activation energy just says how fast it takes.
So, it's important that everything is related,
but it's important to know the boundaries.
Okay. And then, just to come kind of full circle, right,
so these are the reaction energy diagrams the way like,
the way we're used to seeing them,
and then to follow through, right, if we have,
here's our reaction rate, okay, and so,
if we don't have a catalyst, we have our forward
and our reverse reactions.
Okay, and we get to equilibrium, right.
Here's time.
Okay, then, if we add the catalyst, right,
with the catalyst, then, here we go, notice we are getting
to the same place, right.
We just, get, right, so, with catalyst,
equilibrium is reached faster.
Okay. So, at this point,
we've pretty much covered everything we need to know
about the energetics of reactions, so,
this will be a good time to review these notes;
make sure you feel really solid on the concepts.
Find a few practice problems in your textbook
and be ready for class.
I look forward to seeing you.