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MCAT
Course: MCAT > Unit 5
Lesson 11: Principles of bioenergeticsThermodynamics vs kinetics
Created by Jasmine Rana.
Want to join the conversation?
- At, shouldn't it be "transition state", rather than "intermediate"? I was under the impression that whenever there was an intermediate, it would be a local minimum. 3:26(37 votes)
- Correct. This video is really sloppy in terms of language.(39 votes)
- what is the differences between thermodynamics and kinetics?(16 votes)
- Thermodynamics, simply defined, is the branch of physics that deals with the conversion of different forms of energy, and the relations between heat and various energy forms such as mechanical, electrical, or chemical energy.
Kinetics deals more with the actions, or forces, that cause various motions (also known as dynamics). It is concerned with the rate of reactions.(31 votes)
- So I guess you can say an Airplane is the enzyme in your venture to getting to your special someone. Without the enzyme, you're travelling across the world with only a row boat.(15 votes)
- atwhats rxn stand for? 1:13(0 votes)
- Rxn is a shorthand way of writing "reaction."(17 votes)
- Wouldn't the reaction on the left (the spontaneous one) be more kinetically favorable then the reaction on the right (non-spontaneous one)?
You couldn't have a reaction that is Kinetically favorable but thermodynamically unfavorable?(3 votes)- You're correct about the reaction on the left being both spontaneous/thermodynamically favored (deltaG<0) AND kinetically favored (activation E)fwd < (activation E)reverse.
It is possible to have a product that is thermodynamically unfavored but kinetically favored. But I think if you're talking exclusively about a forward vs reverse reaction (like A<=>B), because the transition state is the same energetically for the forward and reverse, the thermodynamically favored reaction will always be kinetically favored. This is because the higher E/less stable reactant is energetically closer to the transition state than the lower E/more stable product.(4 votes)
- To clarify:
Forward Reaction: A -> B
- Thermodynamically favourable (G<0 then spontaneous)
- Kinetically favourable (Assuming Ea is provided from the environment or ignition source. Small Ea)
Backward Reactions: B -> A
- Thermodynamically unfavourable (G>0 then non-spontaneous)
- Kinetically unfavourable (Larger Ea)
If this reaction is an equilibrium reaction, how is the backwards reaction even possible? Sure, the forwards reaction will decrease or deaccelerate until the quotient reaction equal the equilibrium constant. But still, no B can be formed if it is thermodynamically and kinetically unfavourable.(4 votes) - Are all thermodynamically favoured reactions spontaneous? That is, do all thermodynamically favoured reactions have reactants high in Gibbs and products low in Gibbs?(1 vote)
- Thermodynamically favored reactions do have a negative delta G (high free energy in reactants and lower free energy in the products) and thus can be called spontaneous! However, that doesn't mean that every spontaneous reaction will happen immediately!
Think about it like this--burning a log is thermodynamically favorable, however, you need an input of energy for the process to begin (IE: lighting a match). Logs don't randomly catch on fire! (thank goodness!)(4 votes)
- So does this mean that the reaction on the left is both thermodynamically favored (delta G is zero, thus spontaneous) AND kinetically favorable (activation energy is lower), and that the reaction on the right is both thermodynamically unfavored and kinetically unfavored?(2 votes)
- 1) What are the differences between energetic stability and thermodynamics?
2) What does it mean by kinetically stable ?
3) Is activation energy needed to overcome the kinetic energy barrier OR is activation energy the kinetic energy barrier itself?(1 vote) - How does the mixing of a solution affect thermodynamic and kinetic stability?(1 vote)
Video transcript
In this video, I want
to go ahead and talk about thermodynamics
and kinetics. I remember that when I was first
learning about these two things in the context of
chemical reactions, I used to hear phrases like,
oh, this chemical reaction-- let's say hypothetically
A going to B-- is kinetically favorable,
so it must proceed. But then the next day
when I go to class, I'd hear another statement that
said, oh, well this reaction A going to B is
thermodynamically favorable, so it must proceed. So in this video, I want to
go ahead and clarify these two statements for you
and really understand what it means to
thermodynamically favored and/or kinetically favored. So to get it started, I've
already drawn two plots so that we can plot out the
free energy change that occurs for the forward
reaction-- we'll do that on the left side-- and
the reverse reaction, which we'll do on
the right side. So before we get
into that, let's go ahead and label our axes. Our y-axis in both
cases will be measuring free energy, which is
in units of joules. Then our x-axis will be
an abstract dimension called the reaction coordinate,
which essentially allows us to monitor the
progress of a reaction. So now I'm going
to go ahead and say that the forward reaction we'll
draw out in this teal color. And the reverse reaction
we'll go ahead and draw out in this pink color. So let's go ahead and start
with the forward reaction. So I'm just going
to go ahead and say that the forward reaction
has a negative delta G value. So remember that means
it's spontaneous, and visually that means
that our reactants start off at a higher free energy
than our products. Now, this of course means that
the reverse reaction, which will have the same
magnitude of delta G, that is the free energy change, will
be at the same numerical value. But of course,
since the reaction is going the opposite
direction, the sign of delta G will be now positive. And visually, we're
saying essentially that our reactant,
which in this case is B, starts our at a lower free
energy than our product, which is A. Now thus far in drawing
this free energy diagram, we've just been talking
about thermodynamics. But it turns out that there
is also a kinetic energy barrier for the conversion
of reactants to products, regardless of whether
the reaction is spontaneous or non-spontaneous. This kinetic
barrier of energy is referred to as the free
energy of activation, or simply activation energy. So I'm going to go ahead and put
in parentheses E sub A, which we'll say stands for
activation energy. And remember that
delta G of course is talking about thermodynamics. With that said, let's go ahead
and add this kinetic energy barrier to our diagrams. And we can do this
by understanding that the activation energy is
defined as the amount of energy that is required to form
a high-energy intermediate during the course
of the reactions. In other words, in our
hypothetical reaction of A going to B, it proceeds
through an intermediate, that is a high-energy
chemical product that won't last very long, but is
important in the conversion of A to B or vice versa. So in our diagrams
here, we can go ahead and indicate that there
is some intermediate, so in between our reactants
and our products, that is at much higher energy
than everything else. And we can go ahead and
then connect the dots. So go ahead and essentially
draw a line from reactants to products that includes
our high-energy intermediate. And this ultimately
allows us to see the presence of the
activation energy as well as the change
in free energy. So let's go ahead and
actually label these things. So on the left side
here, remember our change in free energy, which is
looking at our reactants compared to our products--
we're ignoring this little bump. The change in free
energy, which I'm going to indicate
with the green line, extends between
these two points. And the same on the right
side, just again extending between the start and
endpoints of our reactions. So I'll go ahead and
indicate that this green line in both cases refers
to our delta G values. And in a different
color, let's say red, I'm going to go ahead and
indicate the activation energy, which takes
into account the change in energy between the
high-energy intermediate and the reactants. So in the case on
the left, that change is indicated here with red. And on the right
side, the change between the intermediate
and the reactant is a bit longer, so we'll go
ahead and indicate that here. Now, activation energy
is an important quantity to take into account, because
in order for molecules to react, they must have enough energy to
overcome this activation energy barrier. Essentially, in the case
of a spontaneous reaction for example, I think of
it like the energy one needs to get a ball to
start rolling down a hill. We all know that
gravity will make a ball roll down a hill, which is
like a negative delta G value, it's telling us
that the reaction is very thermodynamically
favorable. But we need to sometimes
give the ball push in order for the reaction to occur. And so that's kind
of this little help that it needs to go over
before it can actually proceed. For a non-spontaneous
reaction, the idea is essentially the same. We still need to have
some activation energy. But in addition, because it
requires an input of energy, we can think about it as
rolling ball up a hill instead of down a hill. Now in general, the idea is
that the lower this free energy change, the faster a
reaction will occur. And remember I'm saying faster,
so I'm talking about kinetics. I'm talking about the
rate of a reaction. So just to write that out, the
activation energy, the smaller it is, the faster the
reaction will proceed. Now in biochemistry
in particular, it's really important to
distinguish between these two terms of thermodynamics
and kinetics, which we've drawn
out in our diagrams as the change in delta G
over the change in activation energy. Because many biochemical
reactions in our body are kinetically
unfavorable, that is to say they have a very
high energy of activation even if they are
thermodynamically favorable. This is why our bodies have
enzymes, which essentially lower the activation
energy of a reaction. So I went ahead and
drew a dotted white line that's a little
bit lower, so you can see that when an
enzyme is present, the height of the
barrier has decreased. And if it's decreased, the
reaction will proceed faster. Now, there's one analogy that
my chemistry professor used to tell us all the time that
really helped me understand the interplay between
kinetics and thermodynamics as they apply to whether or
not a reaction will occur. So I'm going to go ahead and
scroll down so we can briefly talk about this
analogy, which is I think a fun way to
think about all of this. So let's say you went
to a dating website, because you were looking
for your perfect match. And this dating website told
you that your perfect match lived halfway across the globe. And this is such
a perfect match, and they have all
these algorithms. And if this match were
a chemical equation, we would say that you had a
very, very negative delta G value. That means you would be
very, very spontaneous. But you live halfway
around the world, so in terms of
chemistry language, we might say that you
are kinetically limited. That is to say, you don't have
any way of actually travelling halfway across the world to
meet your special someone. So we might say that you have
a very high activation energy. So from this discussion,
perhaps the biggest take-away is that neither kinetics
nor thermodynamics solely determines whether a
reaction will proceed. It's important to take
both into account. Now one of the
applications of this, in biochemistry
especially, is remember that enzymes lower
the activation energy of a reaction. So we can think of
enzymes as light switches. They can regulate whether a
reaction will proceed or won't proceed. But, of course, a
light switch only works if the light
bulb itself is working. And so the metaphor for
the working light bulb is saying that a reaction
has a negative delta G value. Or if light bulb is
not working, we're saying that the delta
G value is positive. That is, it requires a
new battery or energy for the reaction to proceed.