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Collision theory - Chemical Kinetics Video Lecture | Chemistry for EmSAT Achieve
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FAQs on Collision theory - Chemical Kinetics Video Lecture - Chemistry for EmSAT Achieve
| 1. What is collision theory in chemical kinetics? |  |
Ans. Collision theory is a concept in chemical kinetics that explains how chemical reactions occur at the molecular level. According to this theory, for a reaction to occur, reactant molecules must collide with each other with sufficient energy and proper orientation.
| 2. How does collision theory explain the rate of chemical reactions? | |
Ans. Collision theory states that the rate of a chemical reaction is directly proportional to the number of effective collisions between reactant molecules. Effective collisions are those that occur with sufficient energy and proper orientation to break and form chemical bonds.
| 3. How does temperature affect collision theory? | |
Ans. Temperature plays a crucial role in collision theory. As temperature increases, the kinetic energy of molecules also increases, leading to more frequent and energetic collisions. This increases the likelihood of successful collisions, resulting in a higher reaction rate.
| 4. Does collision theory apply to all types of chemical reactions? | |
Ans. Collision theory primarily applies to reactions involving gas-phase reactants. In these reactions, the reactant molecules are in constant motion and have a higher chance of colliding. However, collision theory can also be applied to reactions involving liquid or solid reactants, although the molecular collisions are less frequent.
| 5. Can collision theory explain the selectivity of chemical reactions? | |
Ans. No, collision theory does not provide an explanation for the selectivity of chemical reactions. While collision theory explains the rate at which reactions occur, the selectivity of reactions depends on factors such as the nature of the reactants, the presence of catalysts, and the stability of the reaction intermediates.
Text Transcript from Video
- To think about collision theory,
let's consider the following reaction.
Here we have atom A reacting
with a diatomic molecule B C
to form a new diatomic molecule A B and C.
According to collision theory,
molecules must collide to react.
So for this example,
atom A has to collide with molecule B C,
in order for the reaction to occur.
Next, the collisions must have
the correct orientation in space
to be an effective collision.
For example, let's say for this reaction,
we have our molecule B-C.
Our molecule B-C approaches
A in this orientation,
and since we're forming
a bond between A and B,
let's say this is the proper orientation.
So this is the way our
collision has to occur
in order for the reaction to occur.
If the diatomic molecule B-C approaches
in the opposite direction,
so let's say we have our atom A here,
and then we have C-B, so the
atom C approaches the atom A,
here this is not the proper orientation
for the reaction to occur.
So this would be no.
So, there has to be a collision,
but the collision has to be
in the proper orientation.
And finally, collisions
must have enough energy.
So, if the collision
doesn't have enough energy,
the molecules, or in this case
the atom and the molecule,
will just bounce off of each other.
If you do have enough energy,
the colliding molecules
will vibrate strongly enough
to break bonds.
So let's go ahead and draw this in.
We're starting with a certain
energy for our reactants.
So right here we're going to draw in
the energy for our reactants.
So we have our atom A,
and we have our molecule
B-C at this point,
and let's say our total energy is
20 kilojoules per mol.
When the atom and the molecule collides,
they need enough energy.
They need enough energy to break this bond
between B and C.
So we're trying to
break this bond in here.
And so, we can find that
energy on our diagram here.
So we're starting with
20 kilojoules per mol,
and we need to get up to here, to 60.
So this is how much energy we need
for the reaction to occur.
And we call this the activation energy,
which is symbolized by E sub-a here.
So this is the activation energy.
And the activation energy is important,
because this is the
minimum amount of energy
that's required to initiate
a chemical reaction,
and for this reaction, we
can see we need to get to
60 kilojoules per mol.
So this point right here is
at 60 kilojoules per mol.
We're starting out with 20.
So 60 minus 20 would, of course, be 40.
So the activation energy
for this reaction,
so Ea, according to our diagram,
is positive 40 kilojoules per mol.
So the energy of the collision must be
greater than or equal to
the activation energy,
and at the top right here,
we're going to get a
transitional structure.
So let me go ahead and draw in
a possible transitional
structure for this reaction.
So we have a bond forming between A and B.
At the same time, we have a bond breaking
between B and C,
and we call this transitional
structure right here,
we call this the transition state.
So our structure right here is called
the transition state.
You might also see this
called the activated complex.
So, the transition state
or the activated complex.
And you can see I've drawn
in partial bonds here.
So the bond between B and C is breaking.
At the same time, we have the
bond forming between A and B.
And so, let's think about an analogy here.
Let's say we have a hill.
So here's my hill right here,
and if we have a ball.
Let's say we have a ball right
here at this end of the hill.
Well, it takes energy to
push the ball up the hill.
And let's say we have enough energy
to get the ball to right here.
Well, in that case, it's
not enough for the ball
to roll down the other side of the hill.
Here the ball's going to roll
back to the starting position.
So that's like thinking about
having not enough energy
and the molecules just
bouncing off of each other.
But if you have enough energy,
so if you're starting out
with the ball right here,
and you have enough energy
to bring the ball to the top of the hill,
so just barely to the top here,
the ball can now roll down.
And so, the ball's going
to end up at the bottom
of the hill right here.
And that's thinking about
formation of your products.
So for our example, our products would be
our new diatomic molecule A-B.
So let me draw that in here.
So here we have A-B, and we
also have plus C at this point.
So this would be our products,
and this represents the
energy of our products.
So let's go here to find
that on our diagram.
So we have our products here.
So what energy is that?
So we go over to here,
and we find the energy,
and let's say that's at 10.
Let's say this is 10 right here.
So the energy of our products
is equal to 10 kilojoules per mol,
and the energy of our
products for this example
is lower than the energy of our reactants.
We started out with 20 kilojoules per mol,
and we ended up with
10 kilojoules per mol.
So to find that difference in energy,
to find that change in energy,
that would be the energy of the products
minus the energy of the reactants,
and for this example, the
energy of the products
is 10 kilojoules per mol, so we have 10,
minus the energy of the reactants.
We started out with 20 kilojoules per mol.
So 10 minus 20 gives us a change in energy
equal to negative 10 kilojoules per mol.
So on our diagram,
we can see this change
in energy right here.
So it's the change in energy.
This right here is a
negative change in energy.
So I'm writing delta E
is equal to, is negative.
So this is an exothermic reaction.
So heat is given off here.
Remember, it takes energy to break bonds,
and energy is given off when bonds form,
and for this reaction,
we're giving off heat.
So this represents an energy diagram
for an exothermic reaction.
Now let's think about another reaction.
So down here we have reaction progress.
So we're starting out with
the energy of our reactants right here.
So this is the energy of our reactants
at 20 kilojoules per mol, let's say.
And we know that to get
to this point right here,
this represents the energy
of the activated complex.
So right there will be
our activated complex,
and that's at 80 kilojoules per mol.
So this difference in energy,
this is our activation energy.
So this is Ea.
That represents our activation energy,
which for this reaction
would be 80 minus 20,
which would be equal to
60 kilojoules per mol.
So we get to here.
We get to the transition state
or the activated complex,
and then this would represent
the energy of our products.
So this our energy of our products here,
which for this reaction,
you can see the energy of our products
is greater than the
energy of our reactants.
And so, if we go over to
here, let's say it's at 40.
So let's say this level is
at 40 kilojoules per mol.
So what's the change in energy here?
So the change in energy for this reaction,
once again, is the energy of the products
minus the energy of the reactants,
so that would be 40 kilojoules per mol.
So, 40 kilojoules per mol
minus we started out with 20.
So, 40 minus 20 gives
us a change in energy
equal to positive 20 kilojoules per mol.
So, positive 20 kilojoules per mol.
So that would be this
difference right here.
So, our change in energy for
this reaction is positive.
And so, this represents
an endothermic reaction.
So the previous example
was an exothermic reaction,
where heat is given off,
and this is an endothermic reaction.
This represents an endothermic reaction,
where heat is absorbed.
let's consider the following reaction.
Here we have atom A reacting
with a diatomic molecule B C
to form a new diatomic molecule A B and C.
According to collision theory,
molecules must collide to react.
So for this example,
atom A has to collide with molecule B C,
in order for the reaction to occur.
Next, the collisions must have
the correct orientation in space
to be an effective collision.
For example, let's say for this reaction,
we have our molecule B-C.
Our molecule B-C approaches
A in this orientation,
and since we're forming
a bond between A and B,
let's say this is the proper orientation.
So this is the way our
collision has to occur
in order for the reaction to occur.
If the diatomic molecule B-C approaches
in the opposite direction,
so let's say we have our atom A here,
and then we have C-B, so the
atom C approaches the atom A,
here this is not the proper orientation
for the reaction to occur.
So this would be no.
So, there has to be a collision,
but the collision has to be
in the proper orientation.
And finally, collisions
must have enough energy.
So, if the collision
doesn't have enough energy,
the molecules, or in this case
the atom and the molecule,
will just bounce off of each other.
If you do have enough energy,
the colliding molecules
will vibrate strongly enough
to break bonds.
So let's go ahead and draw this in.
We're starting with a certain
energy for our reactants.
So right here we're going to draw in
the energy for our reactants.
So we have our atom A,
and we have our molecule
B-C at this point,
and let's say our total energy is
20 kilojoules per mol.
When the atom and the molecule collides,
they need enough energy.
They need enough energy to break this bond
between B and C.
So we're trying to
break this bond in here.
And so, we can find that
energy on our diagram here.
So we're starting with
20 kilojoules per mol,
and we need to get up to here, to 60.
So this is how much energy we need
for the reaction to occur.
And we call this the activation energy,
which is symbolized by E sub-a here.
So this is the activation energy.
And the activation energy is important,
because this is the
minimum amount of energy
that's required to initiate
a chemical reaction,
and for this reaction, we
can see we need to get to
60 kilojoules per mol.
So this point right here is
at 60 kilojoules per mol.
We're starting out with 20.
So 60 minus 20 would, of course, be 40.
So the activation energy
for this reaction,
so Ea, according to our diagram,
is positive 40 kilojoules per mol.
So the energy of the collision must be
greater than or equal to
the activation energy,
and at the top right here,
we're going to get a
transitional structure.
So let me go ahead and draw in
a possible transitional
structure for this reaction.
So we have a bond forming between A and B.
At the same time, we have a bond breaking
between B and C,
and we call this transitional
structure right here,
we call this the transition state.
So our structure right here is called
the transition state.
You might also see this
called the activated complex.
So, the transition state
or the activated complex.
And you can see I've drawn
in partial bonds here.
So the bond between B and C is breaking.
At the same time, we have the
bond forming between A and B.
And so, let's think about an analogy here.
Let's say we have a hill.
So here's my hill right here,
and if we have a ball.
Let's say we have a ball right
here at this end of the hill.
Well, it takes energy to
push the ball up the hill.
And let's say we have enough energy
to get the ball to right here.
Well, in that case, it's
not enough for the ball
to roll down the other side of the hill.
Here the ball's going to roll
back to the starting position.
So that's like thinking about
having not enough energy
and the molecules just
bouncing off of each other.
But if you have enough energy,
so if you're starting out
with the ball right here,
and you have enough energy
to bring the ball to the top of the hill,
so just barely to the top here,
the ball can now roll down.
And so, the ball's going
to end up at the bottom
of the hill right here.
And that's thinking about
formation of your products.
So for our example, our products would be
our new diatomic molecule A-B.
So let me draw that in here.
So here we have A-B, and we
also have plus C at this point.
So this would be our products,
and this represents the
energy of our products.
So let's go here to find
that on our diagram.
So we have our products here.
So what energy is that?
So we go over to here,
and we find the energy,
and let's say that's at 10.
Let's say this is 10 right here.
So the energy of our products
is equal to 10 kilojoules per mol,
and the energy of our
products for this example
is lower than the energy of our reactants.
We started out with 20 kilojoules per mol,
and we ended up with
10 kilojoules per mol.
So to find that difference in energy,
to find that change in energy,
that would be the energy of the products
minus the energy of the reactants,
and for this example, the
energy of the products
is 10 kilojoules per mol, so we have 10,
minus the energy of the reactants.
We started out with 20 kilojoules per mol.
So 10 minus 20 gives us a change in energy
equal to negative 10 kilojoules per mol.
So on our diagram,
we can see this change
in energy right here.
So it's the change in energy.
This right here is a
negative change in energy.
So I'm writing delta E
is equal to, is negative.
So this is an exothermic reaction.
So heat is given off here.
Remember, it takes energy to break bonds,
and energy is given off when bonds form,
and for this reaction,
we're giving off heat.
So this represents an energy diagram
for an exothermic reaction.
Now let's think about another reaction.
So down here we have reaction progress.
So we're starting out with
the energy of our reactants right here.
So this is the energy of our reactants
at 20 kilojoules per mol, let's say.
And we know that to get
to this point right here,
this represents the energy
of the activated complex.
So right there will be
our activated complex,
and that's at 80 kilojoules per mol.
So this difference in energy,
this is our activation energy.
So this is Ea.
That represents our activation energy,
which for this reaction
would be 80 minus 20,
which would be equal to
60 kilojoules per mol.
So we get to here.
We get to the transition state
or the activated complex,
and then this would represent
the energy of our products.
So this our energy of our products here,
which for this reaction,
you can see the energy of our products
is greater than the
energy of our reactants.
And so, if we go over to
here, let's say it's at 40.
So let's say this level is
at 40 kilojoules per mol.
So what's the change in energy here?
So the change in energy for this reaction,
once again, is the energy of the products
minus the energy of the reactants,
so that would be 40 kilojoules per mol.
So, 40 kilojoules per mol
minus we started out with 20.
So, 40 minus 20 gives
us a change in energy
equal to positive 20 kilojoules per mol.
So, positive 20 kilojoules per mol.
So that would be this
difference right here.
So, our change in energy for
this reaction is positive.
And so, this represents
an endothermic reaction.
So the previous example
was an exothermic reaction,
where heat is given off,
and this is an endothermic reaction.
This represents an endothermic reaction,
where heat is absorbed.
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